NOTES 


ON 


QUALITATIVE    ANALYSIS 


BY 

HORACE   G.    BYERS 
1 

PROFESSOR   OF   CHEMISTRY,   UNIVERSITY   OF   WASHINGTON 
AND 

HENRY   G.    KNIGHT 

DIRECTOR   OF   EXPERIMENT   STATION 
UNIVERSITY   OF    WYOMING 


NEW   YORK 
D.    VAN    NOSTRAND   COMPANY 

23  MURRAY  AND  27  WARREN  STREETS 
1912 


\ 


COPYRIGHT,  1912,  BY 
D.  VAN   NOSTRAND   COMPANY. 


Nottoooti 

J.  8.  Cushing  Co.  —  Berwick  &  Smith  Co. 
Norwood,  Mass.,  U.S.A. 


PREFACE 

THE  publication  of  a  new  textbook  on  qualitative  analysis 
requires  an  apology  on  the  part  of  the  authors,  in  view  of  the 
number  and  variety  of  those  already  published.  We  submit 
the  following :  — 

Most  of  the  small  texts  are  designed,  or  at  least  serve,  merely 
as  a  guide  to  the  laboratory  practice  of  the  art  of  analysis. 
Many  ignore  wholly  any  relation  between  the  laboratory  work 
and  the  principles  of  general  chemistry.  The  larger  works 
are  of  two  types :  those  which  are  extremely  detailed  labora- 
tory guides  and  those  which  couple  the  theories  and  laws  of 
chemistry  with  directions  for  analytical  procedure  to  an  extent 
which  makes  them  too  involved  for  use  in  the  ordinary  labora- 
tory course,  where  a  limited  amount  of  time  can  be  devoted  to 
the  subject.  In  addition  there  are  always  special  conditions 
which  confront  the  teachers  in  their  various  environments  which 
render  a  particular  arrangement  desirable. 

In  our  courses  we  present  qualitative  analysis  as  a  part  of  the 
laboratory  work  in  general  chemistry,  and  it  therefore  serves  as 
the  experimental  illustration  of  the  chemistry  of  the  metals.  We 
also  offer  an  advanced  course  which  serves  to  develop  the  art  of 
analysis  and  also  deals  in  a  broader  way  with  the  rationale  of  the 
operations.  We  have  therefore  prepared  these  notes,  which 
represent  the  kind  of  course  we  desire  to  give  our  own  students, 
and  offer  them  for  publication  chiefly  for  home  consumption. 
If  others  have  our  viewpoint,  they  may  find  them  useful. 

We  have  tried  to  keep  in  mind  that  qualitative  analysis  should 
develop  laboratory  technique  and  enable  the  student  success- 
fully to  carry  out  the  identification  of  inorganic  materials. 
It  should  also  widen  the  knowledge  and  appreciation  of  the 
student  of  the  fundamental  principles  of  chemistry  and  enable 
him  to  make  use  of  the  general  laws  and  theories  as  applied 

iii 


258707 


iv  PREFACE 

to  particular  and  individual  questions.  It  should  not  be  com- 
plete within  itself,  but  should  stimulate  to  wide  reading  and 
the  correlation  of  the  occurrence,  preparation,  and  uses  of  the 
elements  and  their  compounds  with  the  reactions  used  for  their 
separation  and  identification.  These  aims  we  seek  to  accom- 
plish through  the  reactions,  the  exercises,  the  analytical  tables, 
and  the  application  of  general  principles  to  particular  cases, 
with  the  hope  that  the  student  will  emerge  from  the  course,  not 
only  a  fair  analyst,  but  with  a  broadened  knowledge  upon  which 
to  develop  his  further  chemical  education.  In  view  of  the  very 
large  number  of  our  students  who  use  the  course  merely  as  an 
adjunct  to  the  chemistry  of  the  metals,  and  who  therefore  deal 
chiefly  with  Parts  II  and  III,  we  have  also  tried  to  present  an 
elementary  system  which  will  not  demand  the  use  of  platinum, 
and  so  permits  a  great  economy  in  equipment  for  large  classes. 

It  is  recognized  that  in  the  regular  course  of  analysis  of 
substances  many  complications  may  arise  which  will  require 
modifications  of  treatment  not  provided  for  in  these  notes,  but 
it  is  hoped  that  the  matter  presented  will  enable  the  student 
tcT  recognize  and  deal  intelligently  with  them  when  they  appear. 
It  is  advised  that  each  laboratory  library  or  reading  shelf  be 
provided  with  reference  books  instantly  available  to  students. 
These  should  include  such  books  as  Bottger's  Qualitative 
Analyse,  2d  Auflage,  Treadwell  and  Hall's  Analytical  Chem- 
istry, Stieglitz'  Elements  of  Qualitative  Analysis,  and  Smith's 
(Alexander)  General  Chemistry. 

In  preparing  the  material  for  these  notes  free  use  of  informa- 
tion, wherever  available,  has  been  made,  and  it  is  impossible 
to  properly  credit  the  sources.  No  pretense  to  any  special 
research  -is  made  by  the  authors. 

Sincere  thanks  are  due  to  Dr.  R.  E.  Rose  for  his  painstaking 

criticism  of  the  manuscript  and  proof. 

H.  C.  B. 
SEATTLE,  Jan.,  1912  H.  G.  K. 


CONTENTS 
INTRODUCTION 

PAGE 

QUALITATIVE  ANALYSIS i 

BASIS  OF  IDENTIFICATION i 

CONDITIONS  PRODUCING  REACTIONS     .        .        .  .        .        .        2 

PROPERTIES  USED  IN  IDENTIFICATION 3 

(1)  State  of  Aggregation  S  .         .         .        .        .  3 

(2)  Color  x".        ........  .3 

(3)  Odor  x 3 

(4)  Taste       .                          '      X                       .  .3 

(5)  Melting  and  Boiling  Points      .....  .3 

(6)  Spectra   ^ -4 

(7)  Solubility^ 4 

IMPORTANCE  OF  VARIOUS  PHASES  OF  WORK      .        .  5 

PART    I 

CHEMICAL  PRINCIPLES  INVOLVED  IN  QUALITATIVE  ANALYSIS     .        .  7 

DEFINITION  OF  SOLUTION 7 

KINDS  OF  SOLUTION     ....                7 

PHENOMENA  OF  SOLUTION "...  8 

HYDRATES  IN  SOLUTION       ...                9 

HYDRATION  OF  IONS     ...                 ......  9 

OSMOTIC  PRESSURE       ...                 10 

VAN'T  HOFF'S  HYPOTHESIS 13 

FREEZING  POINT  OF  SOLUTIONS 15 

BOILING  POINT  OF  SOLUTIONS 17 

ACIDS,  BASES,  AND  SALTS  IN  SOLUTION 18 

ELECTROLYSIS 19 


vi  CONTENTS 

PAGE 

HYPOTHESIS  OF  ARRHENIUS 21 

PHYSICAL  EQUILIBRIUM 23 

CHEMICAL  EQUILIBRIUM 25 

IONIC  EQUILIBRIUM 27 

SOLUBILITY  PRODUCT 30 

QUALITATIVE  ILLUSTRATION 31 

Group  I 31 

Group  II 32 

Subgroup  A 36 

Subgroup  B 39 

Group  III .         .40 

Hydrolysis 42 

Oxidation 48 

Bead  Tests 50 

Group  IV        ...........  51 

Flame  Tests 52 

Group  V 52 

Supersaturation 55 


PART    II 
METAL  ANALYSIS 

GENERAL  DIRECTIONS  .        .        . 57 

Equations        ...........  57 

Precipitation  and  Filtration 57 

Washing         ...........  58 

Decantation 59 

Evaporation 59 

Amount  of  Sample .59 

Confirmation  of  Tests .59 

Tests  of  Reagents 59 

Notebooks 60 

THE  HYDROCHLORIC  ACID  GROUP 60 

Silver 60 

Mercury 61 

Lead 62 

ANALYSIS  OF  GROUP  I 63 


CONTENTS  Vll 

PAGE 

EXERCISES  OF  GROUP  I  .64 

THE  HYDROGEN  SULPHIDE  GROUP  .  65 

SUBGROUP  A  

Mercury ....                                    "5 

Lead       .         .                 .        -                 67 

Bismuth 

Copper    . 

Cadmium 69 

ANALYSIS  OF  SUBGROUP  A           .  ^9 

SUBGROUP  B 72 

Arsenic 72 

Antimony 75 

Tin 

ANALYSIS  OF  SUBGROUP  B  .                .                        ....  77 

EXERCISES    .  •  • 

THE  AMMONIUM  SULPHIDE  GROUP      ...  -79 

Iron        .         .                          .                 ......  80 

Chromium 

Aluminium      ...  

Manganese ^4 

Zinc *5 

Nickel     .  

Cobalt 86 

ANALYSIS  OF  GROUP  III  .... 

EXERCISES    ...                                                 ....  91 

AMMONIUM  CARBONATE  GROUP   .                        .                ...  91 

Barium    ...  •  .92 

Strontium 92 

Calcium  .  •  -93 

ANALYSIS  OF  GROUP  IV       ..                 94 

EXERCISES 95 

THE  SOLUBLE  GROUP  ....  •  95 

Sodium  ....  96 

Potassium  .  .  9^ 

Ammonium 97 

Magnesium 9^ 

ANALYSIS  OF  GROUP  V .        .  98 


viii  CONTENTS 

PART    III 
ACID   ANALYSIS 

PAGE 

INTRODUCTION      .        .        . 101 

Rules  of  Solubility  of  Salts 101 

Grouping  of  Acids 103 

GROUP  I 104 

Silicic  Acid     .         .         .         .         .         .         .         .         .         .         .104 

Hydrosulphuric  Acid       .         .         .         .         .         .         .         .         .104 

Thiosulphuric  Acid          .         .         .         .         .         .         .         .         .104 

Sulphurous  Acid 104 

Nitrous  Acid 105 

Carbonic  Acid 105 

Hypochlorous  Acid 106 

ANALYSIS  OF  GROUP  I .     106 

EXERCISES  ON  GROUP  I 107 

GROUP  II 107 

Hydrochloric  Acid 107 

Hydrobromic  Acid 108 

Hydriodic  Acid 108 

Hydrocyanic  Acid 109 

Hydroferrocyanic  Acid 109 

Hydroferricyanic  Acid no 

PREPARATION  OF  A  SOLUTION  OF  THE  ANIONS 115 

ANALYSIS  OF  GROUP  II no 

EXERCISES  ON  GROUP  II 112 

GROUP  III 112 

Sulphuric  Acid         .         .         .         .         .         .         .         .         .         .112 

Chromic  Acid          .         .         .         . 112 

Arsenic  Acid 113 

Arsenous  Acid 113 

Phosphoric  Acid 114 

Boric  Acid 114 

Hydrofluoric  Acid   .         .         .  .         .         .         .         .         .114 

Silicic  Acid     .         . 115 

ANALYSIS  OF  GROUP  III -.        .115 

EXERCISES  ON  GROUP  III     .  116 


CONTENTS  ix 

PAGE 

GROUP  IV 117 

Nitric  Acid 117 

Chloric  Acid 118 

Permanganic  Acid 118 

ANALYSIS  OF  GROUP  IV 118 

EXERCISES  ON  GROUP  IV 119 

GROUP  V "  .        .119 

Oxalic  Acid 119 

Tartaric  Acid 120 

Acetic  Acid 120 

Other  Organic  Acids 120 

ANALYSIS  OF  GROUP  V 121 

EXERCISES  ON  GROUP  V  .        . 121 

PART   IV 
SYSTEMATIC  ANALYSIS 

PRELIMINARY  EXAMINATION .122 

Closed  Tube  Test   .         .         .        .         .        .         .         .         .         .123 

Bead  Test 124 

Flame  Test 125 

Charcoal  Test .125 

PREPARATION  OF  THE  SAMPLE .126 

A  Liquid         ...         .  .         .         .         .         .         .         .         .127 

An  Alloy         .         .         .  .         .         .         .         .         .         .         .127 

A  Non-metallic  Substance  .         .         .         -         .         .         .         .128 

A.  For  Metal  Analysis  . 128 

B.  For  Acid  Analysis 129 

TABLES  OF  ANALYSIS 130 

PART   V 
THE   RARE   METALS 

PRELIMINARY  STATEMENT ,        .138 

Flame  Test      ...........     139 

Microcosmic  Bead  Test  .         .         . 139 

Grouping  of  Rare  Elements - .     140 


X  CONTENTS 

PAGE 

GROUP  I  .140 

Thallium 140 

Molybdenum 141 

Tungsten         .  .     142 

Tantalum  and  Niobium 143 

GROUP  II      ....  -145 

Gold       ....                                                           .  145 

Platinum 147 

Palladium ....  148 

Osmium 149 

Iridium 149 

Ruthenium  and  Rhodium 150 

Selenium 150 

Tellurium 151 

Germanium 152 

GROUP  III .  ...     152 

Vanadium        .         .         .         .         .         .         .         .         .         .         .152 

Titanium 153 

Zirconium 155 

Uranium 156 

Beryllium 157 

Thorium          .         .         .         .         .         .         .         .         .         .         .158 

Cerium 159 

Indium  and  Gallium 160 

Lanthanum,  Didymium,  Yttrium,  Scandium,  Erbium,  etc.         .         .160 

GROUP  IV 161 

Radium  .         .         .         .         .         .         .         .         .         .         .         .161 

GROUP  V 161 

Lithium  ............     161 

Rubidium  and  Caesium    ...  .  .162 


APPENDIX 

LIST  OF  APPARATUS 165 

REAGENTS  IN  SOLUTION 166 

REAGENTS,  SOLID 171 


CONTENTS  xi 

PAGE 

TABLE  OF  ATOMIC  WEIGHTS        .  .172 

PERIODIC  SYSTEM         .  •     i?3 

TABLE  OF  SOLUBILITIES       .  •     i?4 

METRIC  SYSTEM  —  CONVERSION  TABLES      .  •     i?5 

DEGREE  OF  IONIZATION  —  ACIDS,  BASES,  AND  SALTS          .        .        .     i?7 


INTRODUCTION 

THE  first  problem  which  confronts  the  analytical  chemist 
is  one  of  identification.  He  must  determine  what  elements  or 
compounds  are  present,  whether  this  alone  be  the  ultimate 
object  of  his  search,  or  whether  he  be  also  concerned  with  rela- 
tive quantities.  These  two  phases  of  analytical  chemistry  are 
usually  considered  separately,  and  are  known  as  Qualitative 
Analysis  and  Quantitative  Analysis.  Yet  neither  is  inde- 
pendent of  the  other.  Even  where  we  only  desire  to  know 
what  substances  are  present,  it  is  impossible  to  ignore  the 
relative  amounts.  When  we  wish  to  determine  quantities 
exactly,  it  is  imperative  to  know  the  constituents  in  order 
properly  to  select  methods  of  separation.  It  follows  that 
qualitative  must  always  precede  quantitative  investigation. 

The  complex  substances  with  which  the  analyst  deals  may 
consist  of  mixtures  or  of  compounds ;  i.e.  of  substances  formed 
by  the  distribution  of  one  kind  among  those  of  other  kinds, 
or  of  substances  formed  by  chemical  union  of  their  constituents. 
The  separation  of  mixtures  into  their  component  parts  may 
at  times  be  accomplished  wholly  by  mechanical  means,  but 
usually,  with  mixtures,  always  with  compounds,  whatever  degree 
of  separation  is  required  for  identification  is  accomplished  by 
chemical  methods. 

The  basis  of  identification  is  mainly  empirical ;  i.e.  we  are 
guided  by  the  observation  of  a  limited  number  of  facts  which 
may  or  may  not  be  dependent  upon  scientific  generalizations. 

It  is  obvious  that  such  identification  is  never  absolute,  be- 
cause it  would  only  be  possible  to  assure  ourselves  of  the  com- 
plete identity  between  a  sample  of  a  known  and  one  of  an 


2  INTRODUCTION 

unknown  substance  by  comparison  of  all  the  properties  of  each. 
We  tacitly  assume,  however,  that  the  number  of  different  sub- 
stances is  finite,  and  that  if  two  samples  of  material  possess 
in  common  a  certain  number  of  distinct  properties,  they  will 
be  alike  in  all.  In  choosing  the  properties  which  we  may 
use  for  purposes  of  such  comparison,  we  must  avoid  those 
which  are  common  to  all  substances  or  to  such  a  large  number 
as  to  be  useless,  also  such  as  are  incidental  or  easily  affected 
by  circumstances,  as  size,  position,  temperature,  electrical  con- 
dition. Such  properties  as  may  be  used  we  call  characteristic. 
Even  in  the  case  of  characteristic  properties,  use  must  be  made 
of  several,  and  with  increase  in  the  number  used,  certainty  is 
rapidly  increased.  For  example,  if  a  substance  is  yellow  and 
hard,  it  is  not  safe  to  assume  that  it  is  sulphur ;  but  if,  in  addi- 
tion, we  find  it  is  also  brittle  and  burns  with  a  blue  flame, 
giving  rise  to  a  distinctive  odor,  it  becomes  practically  certain 
that  the  substance  under  consideration  is  wholly  or  partially 
sulphur. 

Ordinarily,  use  is  made  of  properties  of  state  or  condition, 
such  as  color,  odor,  taste,  and  properties  of  reaction,  or  changes 
which  may  be  brought  about  in  particular  forms  of  matter  and 
which  are  peculiar  to  them  either  because  of  the  change  itself 
or  of  the  properties  so  produced. 

Reactions  are  induced  by  some  change  in  conditions  under 
which  we  inspect  the  material  investigated;  the  most  important 
means  by  which  reactions  are  produced  are  (i)  change  of  tem- 
perature, (2)  change  of  electrical  condition,  (3)  change  of  chemi- 
cal condition  by  contact  with  other  substances.  Of  these  the  last 
is  by  far  the  most  important,  and  is  often  used  in  conjunction 
with  one  or  both  the  others.  Contact  of  substances  is  usually 
most  easily  secured  when  they  are  in  the  liquid  condition,  and 
consequently  such  contact  is  most  frequently  procured  by  bring- 
ing the  materials  into  the  liquid  state  by  solution  or  by  fusion. 


INTRODUCTION  3 

Recognition  of  elements  or  compounds,  therefore,  depends 
upon  physical  properties,  and  usually  upon  such  as  are  produced 
as  the  result  of  certain  physical  or  chemical  changes.  Almost 
always  the  determination  is  based  upon  a  quantitative  differ- 
ence, such  as  degree  of  hardness,  specific  gravity,  solubility, 
etc.  The  properties  most  frequently  used  in  qualitative  analysis 
are : 

1 I )  State  of  Aggregation. 

The  existence  of  a  substance  in  the  solid,  liquid,  or  gaseous 
state  under  certain  conditions  is  frequently  characteristic, 
as  also  whether  amorphous  or  crystalline,  whether  viscous 
or  mobile. 

(2)  Color. 

Often  substances  betray  their  identity  by  color  or  color 
changes ;  e.g.  gold,  copper,  iodine,  oxides  of  nitrogen ; 
likewise,  the  colors  produced  by  chemical  reaction  are 
frequently  characteristic. 

(3)  Odor. 

Substances  may  have,  or  may  be  made  to  produce  by 
reaction,  odors  which  are  decidedly  characteristic,  and 
while  not  capable  of  accurate  description  are  recognized 
if  once  familiar.  Examples  are  the  odor  of  chlorine,  of 
acetic  acid,  of  hydrogen  sulphide. 

(4)  Taste. 

Tastes,  also,  are  incapable  of  definite  measurement,  but 
may  frequently  serve  the  analyst  as  means  of  identification, 
as  in  the  case  of  sugar,  aconite,  etc. 

(5)  Melting  and  Boiling  Points. 

The  transition  temperature  from  solid  to  liquid  or  from 
liquid  to  gaseous  condition  is  a  fixed  point  for  all  pure 
substances,  and  often  serves  the  chemist  not  only  qualita- 
tively but  quantitatively. 


4  INTRODUCTION 

(6)  Spectra. 

The  vapor  of  each  element  when  heated  to  incandescence 
gives  rise  to  light  waves  of  certain  definite  character.  If 
these  light  waves  are  dispersed  by  a  prism,  as  in  the 
ordinary  spectroscope,  they  produce  bands  of  color  which 
are  characteristic  of  the  element  which  produces  them. 
No  two  elements  produce  the  same  set  of  color  bands, 
and  the  position  of  the  bands  on  the  prismatic  scale  is  an 
absolute  means  of  identification.  The  delicacy  of  the 
spectroscope  is  also  amazing.  It  is  asserted  that  by  its 
use  the  presence  of  one  third  of  a  millionth  of  a  milligram 
of  potassium  may  be  detected  with  certainty.  Magnesium 
is  one  of  the  most  difficult  elements  to  detect  by  this  method, 
yet  one  hundred-thousandth  of  a  milligram  may  be  so 
detected.  Notwithstanding  this  extreme  delicacy,  the  in- 
strument is  used  to  but  a  limited  extent  in  elementary 
work.  The  reasons  are,  in  part,  that  the  method  furnishes 
no  clew  to  relative  quantities,  and  indicates  only  the  ele- 
ments present,  not  their  state  of  combination  in  the  material 
investigated.  It  may  also  be  said  that  except  for  a  few 
elements,  the  conditions  for  incandescent  vapor  are  not 
easily  obtained,  and  considerable  expense  is  attached  to 
the  purchase  and  use  of  an  instrument  capable  of  wide 
application.  Small  spectroscopes  are  sometimes  used  to 
assist  in  the  detection  of  alkali  and  alkaline  earth  metals. 

For  a  full  discussion  of  the  instrument  and  its  applica- 
tion, the  student  is  referred  to  Landauer,  Craw,  Lockyer 
and  others. 

(7)  Solubility. 

The  solubility  of  substances,  changes  of  solubility  by 
changes  in  chemical  combinations,  and  the  marked  physi- 
cal properties  of  certain  soluble  and  insoluble  substances 
produced  by  chemical  changes  are  of  such  general  appli- 


INTRODUCTION  5 

cation  that  most  of  the  operations  of  qualitative  analysis 
make  use  of  solubility  changes.  The  most  extensively 
used  means  of  identification  is  the  preparation  of  a  solution 
of  the  substance,  and  then,  by  addition  of  certain  known 
reagents,  produce  the  separation  from  the  solution  of 
certain  definite  compounds,  either  as  gases,  liquids,  or 
solids,  which  have  such  definite  properties  as  leave  no  doubt 
as  to  the  materials  producing  them.  There  are  many 
reasons  why  this  method  is  of  such  general  application. 
It  is  simple,  easy  of  manipulation,  does  not  require  expen- 
sive apparatus,  does  not  require  a  high  development  of 
a  scientific  training,  furnishes  a  rough  idea  of  relative 
quantities,  and,  above  all,  conveys  to  the  student  a  general 
knowledge  of  chemical  relations  which  is  required  for 
quantitative  manipulation. 

It  will  be  seen  at  once  that  to  use  successfully  the  properties 
above  mentioned  for  the  purpose  of  identification  requires 
thorough  acquaintance  with  chemical  substances.  This  acquaint- 
ance is  best  secured  by  a  systematic  classification  of  the  sub- 
stances proposed  for  study  and  a  thorough  drill  in  the  changes 
which  they  are  likely  to  undergo.  In  an  elementary  course, 
the  study  is  usually  limited  to  a  small  number  of  the  best 
known  elements  and  such  of  their  compounds  as  are  most 
frequently  encountered. 

The  subject  divides  itself  naturally  into  two  parts :  the 
detection  of  the  metallic  elements  and  of  the  acid  radicals. 
The  metals  are  found  to  fall  naturally  into  groups,  the  members 
of  which  are  related  by  similarity  of  behavior  under  specific 
conditions.  It  is  obvious  that  such  grouping  is  somewhat 
arbitrary,  and  the  particular  one  used  is  largely  a  question  of 
convenience.  The  grouping  used  in  this  outline  is  the  one 
used  most  commonly,  and  is  presented  in  Part  II.  The  acid 


6  INTRODUCTION 

radicals  and  non-metallic  elements  group  themselves  around 
common  reactions  much  less  satisfactorily  than  do  the  metals, 
and  are  less  susceptible  of  systematic  treatment.  The  outline 
suggested  in  Part  III  is  more  or  less  new,  but  follows  the  gen- 
eral plan  found  to  be  most  satisfactory. 

The  necessary  acquaintance  with  the  behavior  of  elements 
and  radicals  under  definite  conditions  is  secured  by  means  of 
a  series  of  preliminary  reactions,  using  known  materials. 
When  familiarity  with  these  reactions  has  been  gained,  and 
not  till  then,  is  it  possible  to  'proceed  intelligently  to  the  identi- 
fication of  unknown  substances. 

In  analytical  work  the  first  requisite  is  accuracy.  When  an 
analysis  is  completed,  the  analyst  must  have  absolute  confidence 
in  his  results.  As  long  as  a  doubt  remains,  more  evidence 
should  be  secured  until  the  doubt  is  resolved.  An  important 
point  to  be  noted  is  that  the  test  of  skill  is  the  detection  of 
small  quantities  of  substances.  The  second  requisite  is  speed. 
Accuracy  is  never  to  be  sacrificed,  but  the  successful  analyst 
must  reach  certainty  without  undue  sacrifice  of  time. 

The  study  of  qualitative  analysis  involves  more  than  merely 
the  development  of  skill  in  the  identification  of  chemical  indi- 
viduals by  means  of  routine  methods  of  separation.  The  large 
number  of  individual  substances  and  their  variations  demand, 
in  addition  to  the  development  of  manipulative  skill,  a  careful 
study  of  the  rationale  of  the  operations  and  of  the  chemical 
principles  involved  in  the  various  steps.  A  sort  of  syllabus 
of  these  principles  is  presented  in  Part  I. 

The  acquaintanceship  gained  by  practical  experience  with 
substances,  the  classification  of  the  knowledge  of  the  relation- 
ship of  the  substances  studied,  and  the  rationale  of  the  processes 
should  go  hand  in  hand,  and  consequently  in  the  following 
pages  the  attempt  is  made  to  supply  at  least  the  topics  of  a 
course  of  lectures  and  recitations  which  shall  accompany  the 
laboratory  work  and  exercises  furnished  by  Parts  II  and  III. 


PART    I 

CHEMICAL  PRINCIPLES  INVOLVED   IN   QUALITATIVE 

ANALYSIS 

SOLUTION.  Since  the  vast  majority  of  the  operations  of 
analysis  are  carried  out  by  means  of  solution  in  water,  a  con- 
ception of  the  present  state  of  knowledge  of  solutions  is  clearly 
important.  Unfortunately  there  is  no  generally  accepted  defi- 
nition of  solution,  and  any  which  may  be  given  involves  some 
points  more  or  less  in  controversy.  Perhaps  as  satisfactory  as 
any  is  the  statement,  "  A  solution  is  a  homogeneous  mixture  of 
two  or  more  substances."  In  using  the  term  "  mixture  "  above, 
the  occurrence  of  chemical  change  during  the  process  of  forma- 
tion of  a  solution  is  not  necessarily  precluded.  Without  enter- 
ing upon  an  academic  discussion  of  the  question  of  the  nature 
of  the  process  of  solution,  it  may  be  as  well  to  consider  that 
chemistry  recognizes  three  conditions  under  which  two  or  more 
substances  may  exist  together,  viz.  as  mixtures,  compounds, 
and  solutions. 

It  is  clear  that,  by  the  above  definition,  solutions  include  the 
following  mixtures : 

(a)   Solids,  liquids,  and  gases  with  gases. 
(£)   Solids,  liquids,  and  gases  with  liquids. 
(c)    Solids,  liquids,  and  gases  with  solids. 

To  the  analyst,  b  is  of  maximum  importance,  and  in  the  fol- 
lowing discussion  attention  will  be  practically  wholly  confined 
to  it.  Also,  from  the  viewpoint  of  the  definition  it  appears  that 
solution  is  to  be  considered  as  a  mutual  act  on  the  part  of  the 

7 


8  QUALITATIVE  ANALYSIS 

components.  Yet  we  ordinarily  speak  of  the  liquid,  or  that 
component  present  in  the  larger  quantity,  as  these/vent,  and  the 
other  component  as  the  sohite.  According  to  the  orthodox  view 
of  the  subject,  solution  is  not  considered  as  involving  chemical 
change ;  as,  for  example,  solution  of  sugar  or  of  salt  in  water  is 
not  a  chemical  process,  though  chemical  changes  may  subse- 
quently take  place  in  the  substance  dissolved,  or  between  the 
components.  There  are,  however,  certain  phenomena  which 
accompany  the  process  which  are  difficult  to  harmonize  with 
this  view.  In  every  case  of  solution,  heat  changes  and  volume 
changes  occur.  Since  it  happens  that  when  the  volume  of  a 
solution  is  less  than  the  sum  of  the  volume  of  its  components, 
heat  is  usually  evolved,  and  when  the  volume  of  the  solution  is 
greater  than  that  of  the  components,  a  lowering  of  the  tempera- 
ture is  observed,  it  might  perhaps  be  inferred  that  the  energy 
changes  are  due  to  the  volume  changes.  It  is,  however,  true  at 
times  that  a  lowering  of  the  temperature  is  coincident  with  con- 
traction of  volume,  and  an  evolution  of  heat  with  expansion  of 
volume.  The  following  experiment  will  illustrate  the  variability 
of  these  phenomena. 

EXPERIMENT  I.  Mix  in  graduated  cylinders  the  following  substances  and 
note  the  change  of  temperature,  and  after  the  room  temperature  is  regained, 
the  change  in  volume. 

(a)    50  c.c.  alcohol  with  50  c.c.  water  (  +  °,  contraction). 

(£)    50  c.c.  sulphuric  acid  with  50  c.c.  water  (  +  °,  contraction). 

(c)    50  c.c.  acetic  acid  with  100  c.c.  water  (  —  °,  contraction). 

(d}  40  c.c.  alcohol  with  60  c.c.  carbon  disulphide  (  — °,  expansion). 

In  a  general  way  it  may  be  said  that  the  properties  of  solu- 
tions are  not  additive.  A  fuller  discussion  of  this  phase  of  the 
subject  may  be  found  in  Ostwald's  Solutions,  in  which  also 
abundant  references  to  the  original  literature  will  be  found. 

There  are  many  cases  where  the  act  of  solution  is  certainly 
accompanied  by  chemical  change.  Examples  are  found  in  the 


CHEMICAL  PRINCIPLES  9 

solution  of  anhydrides  to  form  acids,  of  basic  oxides  to  form 
bases,  and  the  formation  of  hydrates  by  solution  of  anhydrous 
salts.  There  are  other  cases  of  undoubted  chemical  union 
where  separation  of  the  substance  formed  has  not  yet  been  ac- 
complished, such  as  the  formation  of  sulphurous  and  carbonic 
acids,  ammonium  hydroxide,  etc.  Rupert1  reports  the  separa- 
tion of  hydrates  of  ammonium  hydroxide.  There  are  also  ex- 
amples of  the  formation  of  definite  hydrates  of  such  character 
that  the  composition  varies  with  pressure  or  the  temperature  in 
a  manner  different  from  that  of  ordinary  chemical  compounds. 
These  are  represented  by  the  so-called  constant  boiling  mixtures 
of  hydrochloric  acid,  nitric  acid,  hydrobromic  acid,  etc.  with 
water. 

Jones2  claims  to  have  demonstrated  the  existence  in  solution 
of  complex  hydrates  which  cannot  be  isolated  in  a  pure  condi- 
tion, but  which  are  to  be  considered  as  partially  the  cause  of 
the  failure  of  concentrated  solutions  to  conform  to  the  Ostwald 
dilution  law  (vide  infra).  The  student  will  find  a  brief  ex- 
position of  the  Hydration  Theory  of  Jones  in  his  work  on  Physi- 
cal Chemistry  and  a  resume  of  the  Solvate  Theory  in  the 
American  Chemical  Journal,  Vol.  41,  p.  19. 

Washburn3  has  also  demonstrated  the  existence  of  hydrated 
ions  in  the  case  of  solutions  of  the  chlorides  of  the  alkali  metals 
which  are  themselves  never  hydrated  when  crystallized  from 
solution;  and  McGee4  has  shown  the  same  to  be  true  for  the 
chlorides  of  the  alkaline  earths. 

In  view  of  such  facts,  some  chemists  hold  the  view  that  solu- 
tions are  chemical  compounds  of  a  less  definite  character  than 
those  which  can  be  isolated  in  a  pure  condition.  For  a  more 
extensive  view  of  this  phase  of  the  subject,  the  student  is  re- 

1  Jour.  Am.  Chem.  Soc.,  Vol.  31,  p.  866;  Vol.  32,  pp.  748-749;  June,  1910. 

2  Am.  Chem.Jour.,  Vol.  41,  p.  19.          3  Jour.  Am.  Chem.  Soc.,  Vol.  31,  p.  322. 
4  Dissertation  for  Master's  Degree,  U.  o/  IV.,  1911. 


I0  QUALITATIVE  ANALYSIS 

ferred  to  a  review  of  the  work  of  Werner l  and  to  the  work  of 
Jones  2  and  Kahlenberg.3 

The  only  general  view  regarding  the  condition  of  dissolved 
substances  which  has  gained  extensive  credence  is  that  of  van't 
Hoff.  This  hypothesis  has  been  extremely  fruitful  of  results  in 
systematizing  the  reactions  with  which  the  analyst  deals  and  in 
furnishing  illuminating  viewpoints  concerning  the  reasons  for 
the  results  obtained.  It  will  therefore  be  presented  after  the 
experimental  facts  upon  which  it  rests  are  discussed. 

Osmotic  Pressure 

If  any  soluble  substance  be  put  in  the  bottom  of  a  cylinder  and 
water  placed  upon  it,  gradually  the  solvent  will  take  up  the  sub- 
stance, and,  in  the  course  of  time,  the  solute  will  be  found  uni- 
formly distributed  through  the  solvent.  This  phenomenon  may 
be  ascribed  to  a  chemical  reaction  between  the  solvent  and 
solute,4  or  it  may  be  considered  as  due  to  a  special  kind  of 
force  called  osmotic  force.  If  there  is  placed  between  the 
solvent  and  the  solute  a  partition  of  such  character  that 
the  solvent  can  pass  through  and  the  solution  formed  cannot, 
the  same  action  will  take  place.  The  following  experiments 
illustrate  the  point : 

EXPERIMENT  II.  In  a  narrow  cylinder  place  a  layer  of  chloroform,  then  a  thin 
layer  of  water,  and  on  top  superimpose  a  layer  of  ether.  Stopper  securely,  mark 
the  level  of  the  water  layer,  and  allow  to  stand.  The  ether  will  gradually  pass 
through  the  water  into  the  chloroform,  increasing  the  volume  of  the  lower  layer. 

EXPERIMENT  III.  Three  eggs  of  equal  size  may  be  taken  and,  after  re- 
moval of  the  calcareous  portion  of  the  shell  by  means  of  concentrated  hydro- 
chloric acid,  treated  in  the  following  manner :  One  is  placed  in  pure  water, 
another  in  concentrated  calcium  chloride  solution,  and  the  third  is  kept  for 
comparison.  After  twenty-four  hours  one  will  be  found  largely  distended  by 
the  influx  of  pure  water,  the  other  shriveled  by  extraction  of  water. 

1  Am.  Ckem.Jour.,  Vol.  22,  p.  312.        2  Am.  Chem.  Jour.,  Vol.  41,  p.  19. 

3  Jour.  Phys.  Chem.,  Vol.  5,  p.  339.         4  Kahlenberg,  Jour.  Phys.  Chem.,  Vol.  7,  1907. 


CHEMICAL  PRINCIPLES 


II 


If  the  semipermeable  layer  is  so  constructed  that  it  may  be 
held  in  one  position,  the  solvent  passing  through  it  will  exert  a 
pressure  which  is  capable  of  direct  measurement.  This  pressure 
is  known  as  osmotic  pressure.  Such  an  apparatus,  suitable  for 
qualitative  demonstration,  is  prepared  as  follows : 

EXPERIMENT  IV.  Bore  a  hole  in  a  carrot,  scrape  the  outer  skin,  fill  the 
hole  with  a  saturated  solution  of  sugar,  and  stopper  tightly  with  a  cork  carrying 
a  glass  tube  four  feet  or  more  in  length.  Place  the  whole  in  a  beaker  of  pure 
water,  clamp  in  an  upright  position,  and  mark  the  level  of  the  solution  in  the 
glass  tube.  In  a  few  minutes,  the  level  of  the  liquid  will  be  seen  to  be  rising, 
and  the  height  to  which  it  will  go  will  depend  solely  upon  the  strength  of  the 
carrot  cell  walls. 

Such  an  apparatus  is  clearly  not  capable  of  withstanding  very 
great  pressures,  and  hence  is  not  suitable  for  exact  measurement 
of  osmotic  force.  Pf effer 1  prepared  osmotic  cells  in  such  a 
manner  that  he  was  able  to  measure  accurately  the  quantitative 
value  of  this  tendency  of  a  solvent  to  pass  into  and  dilute  a 
solution.  His  apparatus  consisted  of  a  porous  cup  in  the  walls 
of  which  was  deposited  a  film  of  copper  ferrocyanide.  The  de- 
tails of  the  method  of  preparation  of  the  apparatus  are  given  in 
all  works  on  physical  chemistry  and  are  omitted  here. 

Using  this  variety  of  cell,  Pfeffer  obtained  the  following  re- 
sults for  sugar  solutions : 


PER  CENT  SUGAR  IN 
SOLUTION 

PRESSURE  IN  CM.  OF 
MERCURY 

PRESSURE  PER 
i  PER  CENT 

1.  00 
2.00 

53-2 
101.6 

53-2 
50.8 

2.74 
4.00 

6.00 

i5i-3 
208.2 

307-5 

55-4 
52.1 

5i-3 

1.  00 

53-5 

53-5 

1  Osmotische  Untersuchungen,  Leipzig,  1877. 


12 


QUALITATIVE  ANALYSIS 


It  will  be  noted  that  the  results  in  the  third  column  range 
irregularly  around  a  mean  value  and  indicate  that  the  pressure 
is  proportional  to  the  concentration. 

Pfeffer  also  shows  the  following  results  to  be  obtainable  for 
the  effect  of  temperature  upon  osmotic  pressure,  using  a  I  per 
cent  sugar  solution.  (Ostwald,  Outlines  of  General  Chemistry, 
p.  128.) 


TEMPERATURE 
C. 

TEMPERATURE 
ABSOLUTE 

PRESSURE 

PRESSURE 
CALCULATED 

142 

287.2 

5I.O 

51.2 

15-5 

288.5 

52.1 

52.9 

32.0 

305.0 

54-4 

54.2 

36.0 

309.0 

56.7 

55.8 

From  these  and  similar  results  it  appears  that  the  temperature 
coefficient  of  osmotic  pressure  is  the  same  as  that  for  gases, 
viz.  0.00367. 

The  general  conclusions  drawn  from  the  results  of  Pfeffer 
are  that  the  laws  of  Boyle  and  of  Charles  for  gases  are  equally 
applicable  to  the  variation  of  osmotic  pressure  with  changes  in 
concentration  and  temperature. 

Recently  the  work  of  Morse  and  his  coworkers  has  developed 
a  method  of  determination  of  osmotic  pressure  which  leaves 
very  little  to  be  desired  in  the  way  of  accuracy  and  has  shown 
that  the  osmotic  pressure  of  a  sugar  solution  of  unit  concentra- 
tion is  24.45  atmospheres  at  o°  C.  and  increases  with  the  abso- 
lute temperature.  The  following  table  *  is  instructive  : 

1  Am.  Chem.  Jour.,  Vol.  41,  p.  274. 


CHEMICAL  PRINCIPLES 
OSMOTIC  PRESSURES 


WEIGHT        SERIES 

SERIES 

SERIES 

SERIES 

SERIES 

SERIES 

NORMAL           III 

IV 

V 

VI 

VIII 

VII 

CONCEN- 

PRESSURE 

/PRESSURES 

PRESSURES 

PRESSURES 

PRESSURES 

PRESSURES 

TRATION 

0° 

4°-5° 

10° 

15° 

20° 

25° 

O.I 

2.42 

2.40 

2.44 

2.48 

2.52 

2.56 

0.2 

4-79 

475 

4.82 

4.91 

5.02 

5.10 

o-3 

7.11 

7.07 

7.19 

7-33 

7-45 

7-57 

0.4 

9-35 

9-43 

9.58 

9.78 

9.96 

10.12 

0.5 

11.75 

11.82 

12.00 

12.29 

12.49 

12.73 

0.6 

14.12 

14-43 

14.54 

1486 

15.20 

15.42 

0.7 

16.68 

16.79 

17.09 

17-39 

17.84 

18.02 

0.8 

19.15 

I9-3I 

1975 

20.09 

20.60 

20.73 

0.9 

21.89 

22.15 

22.28 

22.94 

23.31 

23.66 

1.3 

24.45 

24-53 

25.06 

25.42 

26.12 

26.33 

Total  pressure      .     .    131.71 

132.68 

134.75 

137-49 

140.52 

142.24 

Mean  ratio  of  osmotic 

to  gas  pressure  .     . 

1.074 

1.065 

1.O6  1 

1.064 

1.069 

1.064 

Mean  loss  in  rotation 

(per  cent)    .     .     . 

1-73 

1.45 

O.22 

0.10 

0.00 

0.20 

Since,  then,  osmotic  pressure  is  proportional  to  the  concentra- 
tion, it  follows  that  Avogadro's  principle  applies  also  to  this 
phenomenon. 

The  Hypothesis  of  van't  Hoff 

This  remarkable  concordance  of  the  gas  laws  and  those  of 
osmotic  pressure  led  van't  Hoff1  to  the  conclusion  that  dissolved 
substances  exert  the  same  pressure,  in  the  form  of  osmotic 
pressure,  as  they  would  exert,  were  they  gasified  at  the  same 
temperature,  without  change  of  volume.  It  will  be  seen  that 
for  cane  sugar  the  results  of  Morse  (see  table  above)  do  not 
confirm  this  conclusion  exactly,  and  the  following  modification  is 
suggested  by  Morse  2  :  "  Cane  sugar  dissolved  in  water  exerts  an 


1  Zeitschr.f.  Physik.  Chem.,  Vol.  I,  p.  481 ;   1887. 

2  Am.  Chem.  Jour.,  Vol.  34,  p.  93. 


I4  QUALITATIVE  ANALYSIS 

osmotic  pressure  equal  to  that  which  it  would  exert  were  it 
gasified  at  the  same  temperature  and  the  volume  reduced  to  that 
of  the  solvent  in  the  pure  state" 

The  value  of  this  view  lies  in  the  fact  that,  if  it  be  correct 
for  all  substances  with  any  kind  of  semipermeable  membrane, 
we  can  carry  over  to  solution  all  the  knowledge  we  possess  con- 
cerning gases,  so  far  as  these  are  dependent  upon  pressure, 
volume,  and  temperature,  provided  only  that  osmotic  pressure  be 
substituted  for  gaseous  pressure.  An  instructive  example  is  as 
follows : 

The  combined  gas  laws  may  be  formulated  pv  =  RT.  (See 
Walker's  Physical  Chemistry,  p.  28.)  If  now  we  consider  the  val- 
ues for  a  gram  molecule  of  a  gas,  under  standard  conditions,  and 
seek  the  value  of  R,  we  obtain  R  =  ^=  IQ33  x  22488  =  ^^ 

The  value  22,488  is  the  gram  molecular  volume  of  gases  in 
cubic  centimeters  as  determined  by  Morley,  and  1033  the  atmos- 
pheric pressure  in  dynes. 

If  now  we  substitute  osmotic  values  as  obtained  by  Morse1 
for  a  volume  normal  solution  of  cane  sugar,  22.61  atmospheres 
at  o°  C,  R  =  '033x22610  =  85;554 

It  follows  that  measurements  of  osmotic  pressure  should  lead 
to  the  determination  of  molecular  weights  just  as  in  the  case  of 
gases.  If,  however,  the  attempt  is  made  so  to  use  it,  the  diffi- 
culty is  quickly  encountered  that  for  many  substances  results 
are  Obtained  which  are  entirely  at  variance  with  facts  otherwise 
clearly  determined.  For  example,  if  we  dissolve  a  gram  mole- 
cular weight  of  potassium  chloride  (74.5)  in  22.4  liters  of  water, 
the  osmotic  pressure  ought  to  be,  by  the  above  examples,  prac- 
tically one  atmosphere.  The  experiment  shows,  however,  that 
it  is  1.88  atmospheres.  Similar  discrepancy  is  also  observed  in 

1  Am.  Chem.  Jour.,  Vol.  34,  p.  85. 


CHEMICAL  PRINCIPLES  15 

the  case  of  practically  all  salts  and  all  the  ordinary-  acids  and 
bases. 

The  explanation  of  this  apparent  failure  of  the  facts  to  con- 
form to  the  van't  Hoff  hypothesis  will  be  presented  when  some 
similar  classes  of  facts  have  been  discussed.  This  explanation 
is  known  as  the  lonization  Hypothesis  of  Arrhenius  (vide  infra). 
Before  closing  the  discussion  of  osmotic  pressure,  attention 
should  be  directed  to  the  fact  that  it  is  possible  to  regard 
osmotic  force  only  as  a  special  manifestation  of  that  general 
property  of  substances  which  for  want  of  a  better  name  we  call 
chemical  affinity.  An  article  by  Kahlenberg1  presents  the  re- 
sults of  experiments  with  various  membranes  and  solvents  and 
conclusions  wholly  at  variance  with  those  just  discussed. 

Freezing  Point  of  Solutions 

When  a  foreign  substance  is  dissolved  in  a  liquid,  the  freezing 
point  of  the  solution  is  lower  than  that  of  the  pure  solvent,  and 
for  moderate  concentrations  the  lowering  of  the  pressure  is 
proportional  to  the  concentration.  This  fact  was  clearly  demon- 
strated by  Blagden,2  but  his  discovery  was  neglected  and  for- 
gotten for  almost  a  hundred  years  until  rediscovered  by  Riidorff.3 
The  deviations  from  this  law  in  concentrated  solutions  are 
probably  due  to  association  of  molecules  of  the  solute  either 
with  each  other  or  with  molecules  of  the  solvent.  This  phase 
of  the  subject  is  rather  fully  presented  by  Jones  in  the  Solvate 
Theory  resume  referred  to  on  page  9.  In  1882  Raoult,4  work- 
ing with  solvents  other  than  water  and  with  water  solutions  of 
organic  compounds  other  than  salts,  arrived  at  the  generaliza- 
tion that  equimolecular  concentrations  of  different  compounds  in 
the  same  solvent  effect  the  same  depression  of  the  freezing  point 

1  Jour.  Phys.  Chem.,  Vol.  10,  p.  141 ;  {1906. 

2  Phil.  Trans.,  Vol.  78,  p.  277;  1788.  3  Pogg.  Ann.,  Vol.  114,  p.  63. 
4  Compt.  Rend.,  Vol.  94,  p.  1517 ;  Vol.  95,  p.  188  and  p.  1030. 


16  QUALITATIVE  ANALYSIS 

For  example,  one  gram  molecule  per  liter  of  most  substances 
lowers  the  freezing  point  of  acetic  acid  approximately  3-9°C. ; 
of  benzene  5.3°C. ;  of  phenol  7.6°  C. ;  of  water  1.89°  C.,  etc. 
These  numbers  are  known  as  the  molecular  depression  of  the 
freezing  point.  A  table  showing  a  large  number  of  Raoult's 
observations  is  found  in  Ostwald's  Solutions,  pp.  209-212. 

A  practical  application  of  the  principle  to  the  determination 
of  molecular  weights  was  made  by  Beckmann,1  and  the  appa- 
ratus for  this  purpose  is  one  of  the  familiar  objects  in  modern 
laboratory  equipment. 

Two  classes  of  deviations  from  the  law  of  Raoult  were  brought 
out  by  himself  and  others.  Some  substances  caused  a  lowering 
of  the  freezing  point  less  than  the  normal,  and  these  values  are 
usually  half  as  large  as  the  molecular  depression  and  suggest 
at  once  that  molecules  of  the  substances  associate  to  form 
larger  molecules.  The  other  deviation  is  found  most  markedly 
in  the  freezing  points  of  acids,  bases,  and  salts  when  dissolved 
in  water,  though  the  variation  is  not  wholly  confined  to  this 
solvent,  but  occurs  in  solutions  of  the  same  substances  in  liquid 
ammonia,  liquid  sulphur  dioxide,  and  to  a  lesser  degree  in  some 
other  solvents.  When  a  substance,  such  as  common  salt,  is 
dissolved  in  water,  a  normal  solution  freezes  at  —  3.46°  instead 
of  —  1.89°,  and  this  excessive  lowering  of  the  freezing  point  is 
common,  to  a  varying  degree,  to  all  solutions  of  acids,  bases, 
and  salts  in  water.  Van't  Hoff  introduced  for  salt  solutions  a 
coefficient  '  /,'  which  represents  the  number  by  which  the  normal 
lowering  must  be  multiplied  in  order  to  give  the  value  actually 
found.  Thus  in  the  case  cited,  i  —  1.83,  i.e.  1.89  x  1.83  =  3.46. 
By  means  of  almost  numberless  experimental  determinations,  it 
gradually  developed  that  this  factor  't'  varies  within  somewhat 
narrow  limits  for  a  given  salt,  and  in  that  case  increases  with 
increasing  dilution.  The  factor  is  never  more  than  two  for  a  * 

1  Z.eit.f.  Phys.  Chem.,  Vol.  2,  p.  638  and  p.  323. 


CHEMICAL   PRINCIPLES  17 

binary  salt,  acid,  or  base,  and  in  ternary  or  quaternary  com- 
pounds may  rise  to  nearly  three  or  four  respectively.  It  is  obvi- 
ously possible  that  this  abnormal  lowering  of  the  freezing  point 
of  water  by  acids,  bases,  and  salts  may  be  due  to  the  formation 
of  increased  numbers  of  molecules  by  dissociation.  If  this  be 
the  case,  then  a  binary  compound,  such  as  common  salt,  may 
dissociate  into  two  portions ;  a  ternary  into  three,  etc. 

EXPERIMENT  V.  In  a  freezing  point  apparatus,  using  a  Beckman  ther- 
mometer, determine  the  freezing  point  of  pure  water  and  solutions,  in  the  same 
sample  of  water,  of  \  normal  sugar  and  |  and  ^  normal  salt. 

Boiling  Points 

As  early  as  1822  Faraday1  attempted  to  determine  the  in- 
fluence of  dissolved  substances  upon  the  boiling  point  of  water. 
Yet  such  were  the  difficulties  presented  by  variation  in  the 
results  obtained  that  it  was  not  until  1889  that  Beckmann 2 
devised  a  method  by  which  the  determination  could  be  made 
with  sufficient  accuracy  for  the  determination  of  molecular 
weights.  The  most  convenient  as  well  as  the  most  recent 
apparatus  for  the  purpose  is  that  of  Menzies.3  The  influence 
of  nonvolatile  dissolved  substances  upon  the  vapor  pressure  of 
solvents,  upon  which  the  boiling  point  depends,  has  been  shown 
to  be  of  the  same  general  character  as  that  upon  freezing  points, 
viz.  that  equimolecular  concentrations  of  substances  produce 
the  same  increase  in  the  boiling  point  of  a  given  solvent,  the 
molecular  elevation  of  the  boiling  point  being  unlike  for  dif- 
ferent solvents. 

Thus  a  grain  molecule  of  most  substances  dissolved  in  1000 
grams  of  the  solvent  raises  the  boiling  points  of  certain  solvents 
as  follows  : 

1  Ann.  Chem.  Phys.  2,  Vol.  20,  p.  324. 

2  Zeit.f.  Phys.  Chem.  Vol.  4,  p.  532. 

3  Jour.  Am.  Chem.  Soc.,  Vol.  32,  p.  1615;  1910. 


!8  QUALITATIVE  ANALYSIS 

Alcohol 1.15°     Acetone 1.67° 

Ether 2.11°     Chloroform     ....     3.66° 

Water 0.52°      Benzene 2.67° 

However,  if  acids,  bases,  or  salts  are  dissolved  in  water,  the 
elevation  of  the  boiling  point  is  found  to  be  greater  than  0.52° 
per  gram  molecule  per  liter,  and  the  variation  from  this, 
which  may  be  called  the  normal  value,  increases  with  increased 
dilution. 

Acids,  Bases,  and  Salts  in  Solution 

In  the  discussion  of  Osmotic  Pressure,  Freezing  Points,  and 
Boiling  Points,  attention  has  been  specially  directed  to  the 
abnormal  behavior  of  acids,  bases,  and  salts  when  dissolved  in 
water.  The  chemical  behavior  of  these  substances  in  water  is 
equally  unique.  This  peculiar  behavior  is  well  brought  out  by 
the  following  experiments  : 

EXPERIMENT  VI.  Place  a  piece  of  zinc  in  each  of  two  test  tubes,  and  on 
one  pour  concentrated  sulphuric  acid  (free  from  water)  and  on  the  other 
dilute  sulphuric  acid. 

EXPERIMENT  VII.  Grind  together  dry  silver  nitrate  and  dry  potassium 
chromate,  and,  noticing  the  absence  of  result,  add  water  to  the  mixture. 

EXPERIMENT  VIII.  Place  in  a  test  tube  a  mixture  of  dry  copper  nitrate 
and  dry  ammonium  carbonate  and  heat  in  a  Bunsen  flame.  Contrast  this 
complex  reaction  with  that  which  takes  place  between  solutions  of  each  of  the 
same  salts  when  mixed. 

Examples  of  a  similar  kind  may  be  multiplied  indefinitely. 
The  conclusion  is  inevitable,  that  not  only  is  the  chemical 
behavior  of  such  substances  vastly  hastened  by  the  presence 
of  water,  but  that  at  times  the  /'planes  of  cleavage"  of  the 
reactions  are  influenced  by  the  water  present. 

It  may  be  imagined,  of  course,  that  this  "catalytic"  effect  of 
the  water  is  due  merely  to  the  more  intimate  contact  of  the  re- 


CHEMICAL  PRINCIPLES  19 

acting  substances  which  is   secured  by  solution.     This  is  not 
necessarily  true,  as  may  be  illustrated  as  follows : 

EXPERIMENT  IX.  To  a  solution  of  phosphorus  in  carbon  disulphide  add 
a  solution  of  iodine  in  the  same  solvent  and  compare  the  velocity  of  the  re- 
action between  the  phosphorus  and  iodine  with  that  between  a  small  bit  of  dry 
phosphorus  when  dropped  upon  powdered  iodine. 

EXPERIMENT  X.  Compare  the  rate  and  character  of  reaction  between  dry 
sugar  and  concentrated  sulphuric  acid  with  that  of  dilute  solutions  of  the  same 
substances. 

EXPERIMENT  XI.  Compare  the  action  of  dilute  hydrochloric  acid  and  a 
solution  of  dry  hydrochloric  acid  in  toluene  upon  marble  or  zinc. 

It  is  evident  that  in  some  manner  water  exerts  a  potent  influ- 
ence upon  the  chemical  behavior  of  acids,  bases,  and  salts.  This 
influence  may  be  shown  to  be  such  that  the  portions  of  the  com- 
pounds entering  into  reaction  are  related  closely  to  those  which 
are  liberated  from  solutions  of  the  same  substances  when  sub- 
jected to  the  influence  of  the  electric  current. 

Electrolysis 

When  the  terminals  from  a  source  of  electricity  are  connected 
by  means  of  various  materials,  three  classes  of  substances  are 
speedily  distinguished  : 

A.  Those  which  act  as  insulators  and  prevent  the  passage  of 
electricity  and  are  known  as  nonconductors.     The  vast  majority 
of  substances  belong  to  this  class.     Examples  are :  nearly  all 
pure  liquids,  including  water ;  dry  substances,  including  acids, 
salts,  and  bases,  except  when  highly  heated  ;  and  the  nonmetals. 

B.  Those  which  admit  the  passage  of  the  current  and  yet 
remain  essentially  unchanged  during  the  passage  of  the  current 
and  wholly  unchanged  after  the  source  of  current  is  disconnected. 
These  are  for  the  most  part  metals,  and  their  mixtures,  or  alloys, 
and  are  known  as  conductors  of  the  first  class. 

C.  Those  which,  while  they  allow  the  current  to  flow,  are 


20  QUALITATIVE  ANALYSIS 

simultaneously  decomposed  by  the  current  or  cause  decomposi- 
tion of  the  terminals  and  are  thus  changed  in  composition. 
These  are  known  as  conductors  of  the  second  class  and  are  of 
the  greatest  interest  in  this  connection.  The  substances  belong- 
ing to  this  class  are  bases,  salts,  and  acids  highly  heated  or  in 
solution  ;  chiefly  in  water  solutions. 

It  is  not  to  the  present  purpose  to  discuss  the  character  of  the 
changes  taking  place  when  electric  currents  pass  through  con- 
ductors of  the  second  class,  but  only  to  call  attention  to  certain 
features  of  the  process ;  which  is  known  as  electrolysis. 

In  the  first  place  it  is  to  be  observed  that  the  components  of 
an  electrolytic  solution  are  separately  nonconductors. 

EXPERIMENT  XII.  Prepare  an  electrolytic  cell  with  platinum  electrodes 
and  a  direct  current  source  of  approximately  10  volts.  Into  this  cell  introduce 
dry  salt  and  pure  water  successively,  taking  care  to  remove  the  one  completely 
before  the  introduction  of  the  other.  Then  pour  into  the  cell  a  solution  of 
salt  in  water. 

EXPERIMENT  XIII.  Use  sugar  in  place  of  the  salt,  as  in  Experiment  XII, 
and  note  that  no  current  passes. 

EXPERIMENT  XIV.  Use  a  solution  of  dry  hydrochloric  acid  in  dry  toluene 
or  benzene  and  note  also  that  no  current  passes. 

A  second  point  of  importance  to  be  noted  is  that  if  we  make 
the  experiment  in  such  a  way  that  all  of  a  given  solution  is  kept 
between  the  platinum  terminals,  —  electrodes,  —  the  amount  of 
current  which  passes  increases  with  increasing  dilution. 

EXPERIMENT  XV.  Prepare  an  electrolytic  cell  consisting  of  a  rectangular 
glass  vessel  (2x10x20  cm.)  and  place  platinum  terminals  the  whole  length 
of  the  ends  of  the  cell.  Connect  with  source  of  current  and  ammeter,  introduce 
concentrated  hydrochloric  acid  to  a  depth  of  i  cm.,  and  note  the  current  pass- 
ing. Then  add  distilled  water  slowly  with  stirring,  and  note  increased  flow. 

A  third  point  to  be  noted  is  that  during  the  process  of  electrol- 
ysis, the  partition  of  the  dissolved  substances  is  of  the  same 
character  as  that  which  takes  place  in  metathetical  chemical 


CHEMICAL   PRINCIPLES  21 

changes.  For  example,  it  can  easily  be  shown  that  when  a  cur- 
rent is  passed  through  a  solution  of  silver  nitrate,  the  silver  is 
deposited  at  the  negative  electrode  and  the  rest  of  the  molecule 
(radical  NO3)  appears  at  the  positive  electrode,  there  to  interact 
with  the  water,  producing  nitric  acid  and  oxygen.  Similarly 
with  sodium  chloride,  the  parts  which  appear  at  the  positive  and 
negative  poles  are  sodium  and  chlorine  respectively.  These 
portions  of  the  compounds  which  thus  travel  to  the  electrodes 
were,  by  Faraday,  called  'ions? 

The  Hypothesis  of  Arrhenius 

From  the  first  observation  of  the  distinction  between  the  man- 
ner of  conductivity  of  solutions  and  of  metals  which  was  made 
by  GrothusMn  1805,  various  hypotheses  designed  to  account 
for  the  process  have  been  advanced.  These  need  not  be  dis- 
cussed here,  except  to  call  attention  to  that  of  Clausius,2  who 
assumed  that  some  of  the  molecules  of  dissolved  substances 
were  ruptured  by  collision,  and  that  the  electric  current  seized 
upon  these  portions  to  effect  its  transport  across  the  interven- 
ing space.  In  1887,  Arrhenius3  published  the  hypothesis  which 
is  at  present  the  prevailing  one  and  which  correlates  the  facts 
concerning  acids,  bases,  and  salts  in  practically  the  following 
form.  Certain  substances  (acids,  bases,  and  salts)  when  dissolved 
in  water  (and  certain  other  solvents)  are  partially  decomposed 
into  parts  (ions)  bearing  positive  and  negative  electrical  charges 
of  equal  amount.  This  hypothesis,  primarily  designed  to  ac- 
count for  the  process  of  electrolysis,  is  developed  from  the  facts 
as  follows : 

The  so-called  abnormal  osmotic  pressure,  freezing  point  de- 
pression, and  boiling  point  elevation  are  always  obtained  with 
those  substances  which  are  electrolytes  and  which  manifest  in- 

1  Arrhenius's  Electrochemistry,  p.  21.  '2  Pogg.  Ann,,  Vol.  100,  p.  353. 

3  Zeit.  f.  Phys.  Chem.,  Vol.  i,  p.  631. 


22  QUALITATIVE  ANALYSIS 

creased  chemical  activity  when  placed  in  water.  The  assump- 
tion of  dissociation  in  solution  would  account  for  these  facts  by 
reason  of  the  increased  number  of  molecules  formed,  thus  giving 
rise  to  a  greater  number  of  particles  than  correspond  to  molecu- 
larly  equivalent  quantities.  Also  if  the  free  radicals  are  formed 
by  solution,  chemical  reaction  in  solution  may  be  considered  as 
essentially  only  a  realignment  of  the  ions. 

These  phenomena  are  most  marked  in  aqueous  solution,  but 
are  by  no  means  confined  to  it,  since  the  same  classes  of  sub- 
stances produce  similar  effects  in  liquid  ammonia,  liquid  sulphur 
dioxide,  and  other  solvents,  though  in  less  extensive  fashion. 

The  assumption  of  partial  dissociation  is  necessitated  by  the 
fact  that  the  abnormal  results  are  ordinarily  not  so  great  as 
would  be  expected  from  complete  dissociation.  For  example, 
the  van't  Hoff  coefficient  '/'  is  less  than  two  for  such  a  salt  as 
NaCl  in  ordinary  dilutions,  but  increases  with  dilution  in  such  a 
way  as  to  indicate  that  two  is  its  maximum  value  at  great  dilutions. 

The  assumption  that  the  parts  of  the  molecules  are  electrically 
charged  is  in  accord  with  the  view  that  a  source  of  current 
maintains  a  constant  potential  of  a  definite  magnitude  and  that 
the  passage  of  the  current  consists  in  the  discharge  of  the 
electrically  charged  ions  against  these  electrodes.  That  the 
positive  and  negative  charges  must  be  of  equal  amount,  though 
the  number  of  ions  is  not  necessarily  equal,  follows  from  the 
fact  that  solutions  of  electrolytes  are  as  a  whole  electrically 
neutral. 

The  hypothesis  as  stated  seems  at  first  to  be  at  variance  with 
certain  fundamental  principles.  Taking  a  simple  case,  when  a 
substance  such  as  hydrochloric  acid  dissociates,  the  ions  can  be 
nothing  other  than  hydrogen  and  chlorine.  Hydrogen  is  practi- 
cally insoluble  in  water,  while  chlorine  is  colored,  has  a  strong 
odor,  and  reacts  with  water,  yet  the  solution  of  hydrochloric  acid 
in  water  evolves  no  hydrogen  and  is  perfectly  colorless  if  pure. 


CHEMICAL  PRINCIPLES  23 

It  should  be  noted,  however,  that  the  ions  are  not  free  elements, 
even  when  of  elementary  composition,  since  they  carry  charges 
of  electricity  of  great  amount  in  comparison  with  their  size,  and, 
moreover,  are  combined  with  water,  as  has  been  shown  by  the 
work  of  Jones  and  Washburn  previously  cited. 

In  similar  manner,  many  more  or  less  apparent  difficulties  in 
the  matter  of  concordance  of  facts  with  the  hypothesis  may  be 
met  and  the  general  conclusion  arrived  at  that  the  hypothesis 
furnishes  a  very  satisfactory  explanation  of  the  phenomena 
presented  by  electrolytes. 

NOTE.  A  very  clear  and  somewhat  comprehensive  discussion  of  the  behavior  of 
electrolytes  during  electrolysis  and  of  the  chemical  behavior  of  ionic  substances  is  presented 
in  Chaps.  XIX  and  XX  of  Smith's  General  Chemistry. 

At  the  same  time  it  should  be  recognized  that  the  hypothesis 
of  Arrhenius  is  not  the  last  word  upon  the  subject.  There  are 
certain  facts  which  are  extremely  difficult  to  harmonize  with  the 
hypothesis ;  as,  for  example,  the  rapid  reaction  between  certain 
compounds  dissolved  in  pyridine,  which  solutions  are  not  electro- 
lytes. Also  the  different  methods  of  determination  of  the  de- 
gree of  ionization,  which  ought  to  give  reasonably  harmonious 
results,  fail  to  do  so.  A  discussion  of  this  phase  of  the  subject 
is  presented  by  Kahlenberg.1 

Notwithstanding  the  unsatisfactory  features  and  the  fact  that 
a  not  inconsiderable  number  of  chemists  deny  the  validity  of 
the  hypothesis,  the  behavior  of  substances  during  the  processes 
of  qualitative  analysis  is  so  clearly  and  satisfactorily  expressed 
in  terms  of  ions,  that  it  is  deemed  advisable  to  use  it  as  a  tool 
until  a  more  comprehensive  and  more  satisfactory  view  is  de- 
veloped. 

Physical  Equilibrium 

Such  changes  as  the  distribution  of  solute  between  immis 
cible  solvents,  the  solution  of  gases  in  liquids,  and  the  deposition 

1  Jour,  of  Phys.  Chem.,  Vol.  5,  p.  339. 


24  QUALITATIVE  ANALYSIS 

of  solids  from  solutions  are  ordinarily  regarded  as  physical 
changes,  and  these  operations  play  important  roles  in  qualita- 
tive work.  A  useful  formulation  of  the  conditions  upon  which 
equilibrium  resulting  from  these  changes  depends  may  be  de- 
duced as  follows,  using  a  simple  case  as  a  type. 

EXPERIMENT  XIX.  Make  a  saturated  solution  of  iodine  in  dilute  potassium 
iodide  solution  and  to  it  add  some  chloroform.  Shake  vigorously  and  allow 
the  chloroform  to  settle  out.  The  iodine  will  be  found  distributed  between 
the  two  solvents. 

NOTE.  In  using  this  example  of  physical  equilibrium,  it  is  not  necessary  to  lose  sight  of 
the  chemical  nature  of  the  union  between  potassium  iodide  and  iodine.  The  union  be- 
tween chloroform  and  iodine  may  be  of  similar  though  not  identical  nature. 

If  we  represent  by  5  the  concentration  of  iodine  in  potassium 
iodide  solution,  it  is  evident  that  the  rate  of  passage  into  the 
chloroform  will  depend  not  only  upon  this,  but  also  upon  the 
attraction  of  chloroform  for  iodine  which  at  a  given  temperature 
is  constant,  and  may  be  represented  by  K.  The  velocity  of 
this  reaction  is  then  V=  KS.  The  tendency  of  iodine  to  go 
from  chloroform  to  potassium  iodide  may  be  formulated  simi- 
larly, V  =  K' S'  \  and  since  the  opposing  tendencies  are  equal 
when  the  distribution  is  complete,  V=  Vf ,  or  S'K1  =  KS,  or 

r-  i^  r- 

—  =  — ,  and  since  K  and  K'  are  constant,  —  =  K. 
S      K.  S 

This  conclusion  may  be  verbally  stated  as  follows :  Physical 
equilibrium  between  reacting  substances  depends  upon  a  con- 
stant ratio  between  the  concentration  of  the  materials. 

This  statement  is  of  general  application,  and  its  use  leads  to 
the  laws  of  Dalton  and  of  Henry  with  regard  to  the  solubility 
of  gases  in  liquids,  and,  if  we  keep  in  mind  the  fact  that  the 
concentration  of  solids  is  their  density,  to  the  known  facts  con- 
cerning saturated  solutions. 


CHEMICAL  PRINCIPLES  2$ 

Balanced  Actions  and  Chemical  Equilibrium 

Nearly  all  chemical  reactions  are  reversible,  and  in  analytical 
operations  are  more  or  less  incomplete.  As  a  rule,  also,  in  ana- 
lytical operations  it  is  desired  that  reactions  be  made  as  com- 
plete as  possible.  It  is  therefore  essential  that  the  principles  of 
equilibrium  be  thoroughly  understood  if  the  operations  are  to 
be  intelligently  followed.  A  few  examples  of  reversible  reac- 
tions will  be  helpful. 

EXPERIMENT  XVI.  Into  a  eudiometer  filled  with  mercury  and  inverted 
in  a  mercury  trough  introduce  about  10  c.c.  of  dry  gaseous  ammonia.  Pass 
electric  sparks  from  a  Rhumkorf  coil  through  the  gas  until  the  volume  re- 
mains constant  ;  then  by  means  of  a  pipette  introduce  a  few  drops  of  dilute 
sulphuric  acid,  and,  after  noting  the  absorption  of  the  undecomposed  ammo- 
nia, continue  sparking  until  no  further  contraction  takes  place. 

EXPERIMENT  XVII.  To  a  concentrated  solution  of  magnesium  chloride 
add  ammonium  hydroxide  as  long  as  a  precipitate  forms.  To  the  mixture 
now  add  ammonium  chloride  until  the  precipitate  is  dissolved. 

EXPERIMENT  XVIII.  To  400  c.c.  of  water  add  5  c.c.  each  of  decinormal 
solutions  of  ammonium  sulphocyanate  and  ferric  chloride,  and  divide  into 
four  equal  portions.  To  one  add  a  few  cubic  centimeters  of  concentrated  ferric 
chloride  ;  to  another  a  few  cubic  centimeters  of  concentrated  ammonium  sul- 
phocyanate, and  to  the  third  concentrated  ammonium  chloride,  and  compare 
the  colors  of  these  portions  with  that  of  the  fourth  portion. 

In  the  cases  given  we  are  dealing  with  balanced  actions  as 
expressed  by  the  following  : 


II.    MgCl2  +  2NH4OH$:Mg(OH)2-r-2NH4Cl. 
III.    FeCl3  +  3  NH4CNS  ^  Fe(CNS)3  +  3  NH4C1. 

In  I  it  will  be  observed  that  the  decomposition  ceases  before 
all  the  ammonia  is  decomposed,  and  that  when  the  residual 
ammonia  is  removed,  the  reaction  proceeds  in  the  reverse  di- 
rection, and  since  the  ammonia  is  continuously  removed,  the 
reaction  proceeds  to  completion.  In  II  a  little  investigation 


26  QUALITATIVE  ANALYSIS 

would  show  that  precipitation  ceases  before  the  reaction  reaches 
completion,  and  it  is  observed  that  the  addition  of  ammonium 
chloride  drives  the  reaction  in  the  reverse  direction  with  solution 
of  the  magnesium  hydroxide.  In  III  we  may  trace  the  direc- 
tion of  change  and  the  influence  of  the  relative  quantities  of  the 
reacting  materials  by  the  color  change  due  to  varying  quantities 
of  ferric  sulphocyanate. 

It  is  evident  that  in  each  of  the  above  cases  we  are  dealing 
with  reactions  which  can  proceed  in  either  direction,  and  that 
when  under  given  conditions  reaction  ceases,  i.e.  equilibrium  is 
reached,  it  is  because  the  velocities  in  opposite  directions  are 
equal,  and  that,  since  the  point  where  equilibrium  is  reached  is 
changed  by  varying  the  relative  quantities  of  the  reacting  mate- 
rials, the  speed  of  a  reaction  is  influenced  by  the  relative 
amounts  of  the  reacting  substances.  This  principle,  more  or 
less  clearly  recognized  by  Berthollet,  Williamson,  and  others  in 
special  cases,  was  first  formulated  as  a  general  principle  by 
Guldberg  and  Waage  in  I86/,1  and  is  variously  known  as  Guld- 
berg  and  Waage's  Law,  the  Law  of  Mass  Action,  and  the  Law 
of  Chemical  Equilibrium.  It  may  be  stated  as  follows :  The 
rate  of  a  chemical  reaction  is  proportional  to  the  active  mass 
of  each  of  the  reacting  substances,  and  equilibrium  in  a  react- 
ing system  is  reached  when  the  rates  of  reaction  in  opposite 
senses  are  equal. 

A  most  useful  quantitative  expression  of  this  law  may  be 
arrived  at  by  considering  a  general  reaction  expressed  by  the 
formula 

A+B+^C+D, 

where  the  components  form  a  reversible  reaction  and  are  pres- 
ent in  a  homogeneous  mixture,  i.e.  as  gases  or  in  solution. 

At  any  given  constant  temperature  the  speed  of  reaction  of 

1  An  abstract  appears  in  Jour.  Prakt.  Chem.  2,  Vol.  19,  p.  69  (1879) ;  also  Ostwald's 
Klassiker,  No.  104. 


CHEMICAL  PRINCIPLES  27 

A  upon  B  depends  upon  their  nature  (i.e.  upon  the  driving 
force  of  all  reactions,  which  is  ordinarily  called  affinity,  and 
which  is  constant)  and  upon  the  concentrations.  Concentration 
is  ordinarily  expressed  in  terms  of  moles  per  liter.  If,  then,  we 
have  one  mole  each  of  A  and  B  per  liter,  the  velocity  of  the  reac- 
tion (  V)  at  any  given  instant  is  V=  F,  where  F  is  the  affinity 
constant.  If,  however,  the  concentration  of  A  is  clt  and  B  is  £2, 
then  V—  c^  x  c2  x  F.  Similarly,  if  we  consider  C  and  D,  the 
velocity  of  their  reaction  in  the  opposite  direction  is  expressed 
by  V1  =  r3  x  c±  x  F'  y  where  CB  and  c±  represent  the  concentra- 
tions of  C  and  D  respectively  and  F'  the  affinity  constant  of 
C  and  D.  When  equilibrium  ensues,  however,  V  must  equal 

c   x  c       F' 

V1  ',  and  consequently  c±  x  c2  x  Fl  =  cz  x  c±  x  F1,  or  -     •-*=  -~- 

£3  x  c±      r 

C     X  C 

Since,  also,  F'  and  F  are  constant,  we  have  —  -  -  =  K,  where 

^3    X     C4: 

K  is  the  affinity  constant  of  the  reversible  reaction.  This  is 
the  fundamental  law  of  chemical  equilibrium. 

It  is  evident  that  in  the  case  of  reactions  such  as 


we  can  write  as  follows  :  A  +A  -f  B^C  -\-  C+  D,  and  conse- 
quently  F=  £j  x  cl  x  c2  x  F  and  V  =  cs  x  ^3  x  r4  x  F1,  and  the 

final  form  c\^  =  K. 
c*  x  c, 

NOTE.  Fuller  discussions  of  the  subject  of  balanced  actions  may  be  found  in  Smith's 
General  Chemistry,  Chap.  XV,  Walker's  Physical  Chemistry,  Chap.  XXII,  and  Nernst, 
Chap.  I,  Book  III. 

Ionic  Equilibrium 

Having  presented  in  an  elementary  form  the  fundamental 
principles  of  equilibrium,  it  is  now  in  order  to  apply  them  to 
ionization.  Taking  a  simple  case,  acetic  acid,  we  formulate 
the  change  taking  place 


28  QUALITATIVE  ANALYSIS 

This  is  a  reversible  reaction,  since  by  removal  of  the  water  the 
unionized  acid  is  obtained,  and  it  therefore  follows  that  at  a 
given  dilution,  there  must  be  present  not  only  ions  of  hydrogen 
and  of  the  acetic  radical,  but  also  molecules  of  acetic  acid,  and 
these  are  in  equilibrium  according  to 

IT 

It  also  follows  that  the  more  dilute  the  solution,  the  greater  the 
dissociation.  This  is  a  necessary  consequence  if  the  general- 
ization be  true;  for  if  we  have  at  a  given  dilution  100  each  of 
hydrions  and  acetions,  then 

100  x  100      „ 

—  =  A 


where  by  'a'  concentration  of  the  undissociated  molecules  is 
indicated.  If  now  the  solution  be  diluted  to  10  volumes,  if  no 
change  of  ionization  results,  we  have  in  unit  volume 

10XKW. 

a 
10 

In  order  that  equality  remain  —  must  diminish  in  value,  and 

10 

this  can  only  occur  by  further  dissociation,  giving 

Q+  io)Q+  io) 
=  A. 

a  —  x 

10 

We  have  in  conductivity  measurements  the  most  convenient 
means  of  determining  the  degree  of  ionization  at  different  dilu- 
tion, and  consequently  the  value  of  the  ionization  constant,  K. 
An  example  of  the  derivation  of  the  constant  and  a  demonstra- 
tion that  it  is  a  constant  is  furnished  from  the  conductivity 
measurement  for  acetic  acid.  The  conductivity  of  acetic  acid 


CHEMICAL   PRINCIPLES 


at  25°  at  infinite  dilution,  that  is,  when  it  is  completely  dis- 
sociated, is  yux  =  364. 

The  conductivity  at  other  dilutions  (ftt.)  divided  by  ^  gives 
the  percentage  of  ionization,  and  these  values  substituted  in  the 
formula  give  the  values  in  column  four  of  the  subjoined  table. 


V 

M. 

PER  CENT  IONIZATION 

K 

8 

4-34 

I.I9 

O.OOOOlSo 

16 

6.10 

1.673 

0.0000179 

32 

8.65 

2.38 

0.0000182 

64 

12.09 

3-33 

0.0000179 

128 

16.99 

4.68 

0.0000179 

256 

23.82 

6.56 

O.OOOOlSo 

In  making  these  substitutions  in  the  formula 


s 


necessary  to  take  into  account  the  dilution,  e.g. 
0.00119     0.00119 

o 

=  0.0000180 ; 

0.99881 


8 


or  in  general  for  a  binary  electrolyte, 

M*  K 

(i  -M)V 

This  is  known  as  OstwalcTs  dilution  law  and  is  found  to  hold 
rigidly  for  weak  acids  and  bases. 

The  fact  that  when  in  this  manner  the  attempt  is  made  to 
determine  the  value  of  K  for  strong  acids  and  bases  and  for  the 
highly  dissociated  salts,  no  constant  value  is  obtained,  has  as 
yet  received  no  adequate  explanation.  A  fuller  and  a  non- 
mathematical  discussion  of  this  phase  of  the  subject  is  found  in 
Walker's  Physical  Chemistry,  pp.  224-240. 


30  QUALITATIVE  ANALYSIS 

EXERCISE 

The  student  is  asked  to  calculate,  from  the  data  given  in  the 
appendix,  the  values  of  K  for  the  acids,  bases  and,  salts.  This 
table  is  to  be  preserved  and  used  in  the  subsequent  discussion. 

The  Solubility  Product 

A  most  interesting  deduction  from  the  law  of  physical  and 
chemical  equilibrium  concerns  itself  with  the  situation  which  ob- 
tains in  a  saturated  solution  of  an  electrolyte  when  in  the  pres- 
ence of  the  solid  substance. 

The  dissolved   material   consists  of   ions  and    undissociated 

molecules,  according  to  the  formula  —  =  K,  but  clt  the  concen- 
tration of  the  undissociated  substance,  is,  in  the  presence  of  the 
undissolved  substance,  the  same  as  the  value  5  in  the  formula 

o 

— -=  K,  and  since  5',  the  concentration  of  the  solid,  is  a  constant, 
o 

then  c1  is  a  constant  and  c2  x  cz  =  Kc^  or  c2  x  c3  =  K. 

This  latter  value  is  known  as  the  solubility  product,  and  the 
conclusion  is  that  in  saturated  solutions  the  product  of  the  con- 
centrations of  the  ions  is  a  constant.  The  importance  of  this 
conclusion  is  manifest  when  we  consider  that  in  saturated  solu- 
tions the  value  c1  is  unchangeable,  yet  the  total  amount  of  a 
given  substance  in  solution  may  be  materially  affected  by  addi- 
tion of  a  very  soluble  substance  containing  a  common  ion.  Ex- 
tensive application  of  this  principle  is  encountered  in  quantitative 
as  well  as  in  qualitative  analysis.  The  student  will  have  already 
encountered  many  examples  of  the  sort,  and  a  single  experiment 
will  be  sufficient  illustration. 

EXPERIMENT  XX.  Make  a  saturated  solution  of  barium  chloride,  and  after 
filtering,  add  a  few  cubic  centimeters  of  concentrated  hydrochloric  acid.  Note 
the  precipitation  of  BaCl2  and  explain  in  terms  of  the  solubility  product. 


QUALITATIVE    ILLUSTRATION 

The  more  important  applications  of  the  foregoing  principles 
will  be  discussed  in  the  following  paragraphs,  and  certain  addi- 
tional topics  will  be  introduced  as  convenient  illustrative  mate- 
rial is  furnished.  In  this  connection  the  metal  groups  will  be 
considered  in  order  and  chiefly  from  the  standpoint  of  the 
reactions  of  the  metal  ions  which  are  used  in  the  group  separa- 
tions recommended.  The  order  in  which  the  groups  are  con- 
sidered will  be  the  same  as  that  in  which  they  are  taken  up  in 
Part  II. 

Group  I 

The  metals  of  this  grou£  are  silver,  lead,  and  mercury  (ous), 
and  are  considered  together  because  their  salts  in  solution  in 
water  form  insoluble  chlorides  when  treated  with  solutions  con- 
taining chlorine  ions.  The  soluble  salts  furnish  the  positive 
ions  Ag+,  Pb++,  and  Hg+^j  The  reactions  with  chloride  ions 
may  then  be  expressed  : 

Ag+  +    Cl-^AgCl 


+    Cl->HgCl 

These  chlorides  are  the  most  insoluble  of  the  more  usual  salts 
of  mercurous  mercury  and  of  silver,  and  so  these  metals  are 
practically  removed  from  the  solution.  Lead  chloride  is  more 
soluble  than  lead  sulphide,  and  consequently  when  the  filtrate 
from  the  precipitate  is  treated  with  hydrogen  sulphide,  lead  sul- 
phide is  formed.  Lead,  therefore,  belongs  in  both  the  first  and 
second  groups.  The  separation  of  the  lead  chloride  depends 

31 


32  QUALITATIVE  ANALYSIS 

upon  its  rapid  increase  in  solubility  with  rise  of  temperature, 
while  the  other  chlorides  remain  essentially  insoluble.  The  sep- 
aration of  silver  from  mercurous  chloride  depends  upon  the 
presence  in  ammonium  hydroxide  of  ammonia  molecules  accord- 

ing to  the  equilibrium  reaction   NH3  +  H2O  ±£  NH4OH,  and 

+ 
the  formation  of  the  complex  ion  Ag(NH3)2  which  is  soluble  in 

water.     We  have  then  the  complex  reaction 

NH 
AgCl  ±£-Agei  «£  Ag  +  01  +  NH8  +  H20  \\ 


Ag(NH3)2±£Ag(NH3)2Cl 

+ 
Since  the  formation  of  the  complex  ions  withdraws  Ag  ions  from 

solution,  AgCl  must  ionize,  and  consequently  AgCl  goes  into 
solution.  When  to  this  solution  an  acid  is  added,  the  with- 
drawal of  ammonia  reverses  the  reaction,  and  silver  chloride 
reappears. 

Similar  complex  ions  with  ammonia  are  formed  with  copper, 
cadmium,  cobalt,  nickel,  amizinc  ions.  Mercurous  ions  react 
with  ammonium  ions  to  form  an  amido  derivative  NH^Hgj', 

which    probably   subsequently    decomposes    into   the    mixture 

+  --  -  —  -  -  — 

NH2Hg+Hg.     The  following  equations   may  be  used    to  ex- 

press these  changes  : 

NH4OH  +  2  HgCl  ->  NH2Hg2OH  +  2  HC1 
NH2Hg9OH  +  HC1  ->  NH2Hg2Cl  +  H2O 

\ 

NH2HgCl  +  Hg 

Group  II 

The  special  characteristics  made  useful  for  the  separation  of 
Group  II  from  Group  I  and  Groups  III  to  V  inclusive  are  that 
while  the  chlorides  are  soluble  in  the  presence  of  a  small  excess 
of  acid,  the  sulphides  are  insoluble  in  a  solution  of  not  too  great 
acidity. 


QUALITATIVE  ILLUSTRATION  33 

The  ions  which  may  be  formed  by  members  of  the  group  are 
as  follows  : 

POSITIVE  IONS^ 

Hg+  and  Hg++  Cd++ 

Pb++  and  Pb++++  As+++  and  As+++++ 

Cu+  and  Cu++  Sn++  and  Sn++++ 

Sb+++ 


Complex  negative  ions  in  which  the  metals  are  present  : 

Cu  (CN)4          Sb63  :          Sb64  and  SbO3 

Cd  (CN)4          As63  :          As64  and  AsO3 

1  1 
Sn03 

Other  negative  elements  and  radicals  may  take  the  place  of 
the  oxygen  and  (CN)  in  these  ions.  This  is  particularly  true 
of  sulphur. 

The  precipitation  of  these  ions  is  dependent  upon  the  fact 
that  hydrogen  sulphide  reduces  the  complex  negative  ions  to 
the  positive  forms,  and  upon  the  exceedingly  small  solubility 
product  of  the  sulphides., 

(The  subject  of  oxidation  and  reduction  will  be  discussed  in 
connection  with  Group  III  (q.v.)). 

The  sulphides  of  the  group  are  by  no  means  of  equal  insolu- 
bility, and  the  precipitation  is  naturally  in  the  order  of  relative 
insolubility,  if  the  concentration  of  the  ions  of  each  metal  is  the 
same,  for  the  reason  that  if  this  be  the  case  with  a  given  concentra- 
tion of  precipitant,  the  solubility  product  will  be  reached  soon- 
est  for  the  sulphide  which  isjgast  sol  n  hip  This_gpnsideration 
TiTof  importance,  especially  when  taken  in  connection  with  the 
influence  of  an  excess  of  acid,  as  thft  following  ^vperiment  will 
indicate. 

EXPERIMENT  XXI.  Mix  together  in  a  tall  cylinder  equal  volumes  of  normal 
solutions  of  copper  sulphate,  cadmium  sulphate,  and  zinc  sulphate.  Add  dilute 


34  QUALITATIVE  ANALYSIS 

ammonium  sulphide,  drop  by  drop,  allowing  the  precipitate  to  settle  as  fast  as 
formed.  The  sulphides  will  form  in  well-defined  layers  in  the  order  of  the 
solutions  named. 

This  experiment  strikingly  demonstrates  the  order  of  solubil- 
ity of  these  sulphides  and  emphasizes  the  need  of  complete 
precipitation  of  each  group,  since  otherwise  the  more  soluble 
members  might  easily  escape  detection  in  the  proper  place,  and 
by  appearing  in  a  subsequent  group  cause  more  or  less  con- 
fusion. 

The  degree  of  acidity  of  the  solution  is  a  matter  of  impor- 
tance, since  precipitation  of  the  sulphides  can  only  take  place 
when  the  product  of  the  ions  exceeds  the  solubility  product. 
With  a  given  concentration  of  cathion,  the  concentration  of  the 
anion  may  easily  be  depressed  in  the  case  of  a  weak  acid  to  an 
extent  which  amounts  to  practical  elimination,  as  is  clear  from 
the  following  equation.1 

/z2  x  s 

—  =A"=a  very  small  number.     (The   small   letters   are 

v 

used  to  indicate  the  molecular  concentration  of  the  correspond- 
ing ions  or  molecules.) 

If  to  such  a  substance  we  add  a  large  concentration  of  hydro- 
gen ions,  then  —  MJ  ""^)=  ^  ancj  jf  x  js  relatively  large, 

y  +  "2s  : 

y  becomes  very  nearly  equal  to  s.  Such  being  the  case,  it  may 
happen  that  m  xs<K,  and  no  precipitation  will  take  place. 
On  the  other  hand,  if  the  concentration  of  acid  is  unduly  small, 
the  ion  product  of  the  more  insoluble  of  the  sulphides  of  Group 
III  may  be  exceeded  and  these  be  precipitated  with  Group  II. 
A  couple  of  experiments  will  make  these  points  clear. 

1  It  should  be  observed  that  hydrogen  sulphide  ionizes  primarily  into  H  and  HS  ions 
but  the  hydrosulphides  decompose  at  once  into  sulphides  and  hydrogen  sulphide,  e.g. 
Cu(SH)2->-CuS  +  H2S.  This  behavior  of  hydrosulphides  is  analogous  to  the  decomposi- 
tion of  hydroxides,  but  is  ordinarily  more  spontaneous.  It  is  therefore  somewhat  simpler 
to  formulate  the  reactions  as  if  divalent  sulphur  ions  entered  directly  into  them. 


QUALITATIVE  ILLUSTRATION  35 

EXPERIMENT  XXII.  Prepare  a  normal  solution  of  cadmium  sulphate  and 
divide  into  two  portions.  To  one  portion  add  a  few  cubic  centimeters  of 
dilute  acid ;  to  the  other  some  concentrated  hydrochloric  acid,  and  pass  hy- 
drogen sulphide  into  both.  In  the  one  case,  observe  the  copious  precipitation, 
and  in  the  other  the  nonformation  of  the  sulphide. 

EXPERIMENT  XXIII.  Prepare  a  normal  solution  of  ferrous  sulphate  and 
divide  into  two  portions:.  To  one  add  a  few  cubic  centimeters  of  dilute  hy- 
drochloric acid  and  pass  hydrogen  sulphide  into  each.  Filter  the  '  neutral ' 
solution  and  add  a  few  crystals  of  sodium  acetate  and  note  the  precipitation 
of  more  ferrous  sulphide.  (Explain  the  action  of  the  acetate.) 

From  these  experiments,  it  is  clear  that  the  solution  for  pre- 
cipitation of  Group  II  must  be  neither  too  acid  nor  too  nearly 
neutral.  Experience  has  shown  that  a  convenient  concentration 
is  about  ^  normal.  A  modification  of  this  statement  is  neces- 
sary in  case  arsenates  are  present.  (See  discussion  of  sub- 
group B.) 

The  separation  of  sub-groups  A  and  B  depends  upon  the  am- 
photeric  nature  of  arsenic,  antimony,  and  tin.  Arsenic  trioxide 
is  soluble  in  both  acids  and  bases,  more  easily  in  the  latter.  It 
is  therefore  both  acid  and  basic  in  character.  The  sulphides 
also  show  the  same  dual  nature,  though  somewhat  less  markedly. 
As  is  usually  the  case,  however,  an  increased  valency  is  accom- 
panied by  increased  acidic  properties,  and  arsenic  pentoxide 
wholly  fails  to  react  with  acids,  and  the  pentachloride  rapidly 
hydrolyzes  in  water.  Similarly  the  pentasulphide  is  less  stable 
than  the  trisulphide,  but  reacts  more  readily  with  bases.  How- 
ever, the  fact  that  the  sulphides  of  arsenic  fail  to  dissolve  readily 
in  acids  is  probably  due  rather  to  their  very  great  insolubility 
in  water  than  to  their  acid  character.  In  solutions  of  alkaline 
sulphides  the  following  reactions  take  place,  with  the  formation 
of  soluble  salts : 

As2S3+  3(NH4)2S^2(NH,)3AsS3 
As2S6  +  3(NH4)2S-*2(NH4)3AsS4 


36  QUALITATIVE  ANALYSIS 

These  reactions  are  analogous  to  the  formation  of  salts  by 
the  union  of  anhydrides  and  basic  oxides  as 

As205  +  3K20->2K3As04 
SO3  +  K2O->K2SO4 

On  account  of  the  stronger  acidic  character  of  As2S5,  its  solu- 
tion is  more  rapid  than  that  of  As2S3.  If,  therefore,  yellow 
ammonium  sulphide,  (NH^S*,  is  used  with  the  trisulphide,  the 
excess  sulphur  reacts 

As2S3  +  2  S->As2S5 

and  the  solution  is  facilitated.  This  last  consideration  is  of 
special  importance,  since  Sb2S3  and  SnS2  are  almost  insoluble 
in  colorless  ammonium  sulphide  (NH4)2S,  but  by  the  yellow  sul- 
phide are  changed  to  Sb2S5  and  SnS2  and  then  dissolve  to  form 
the  soluble  salts 

Sb2S5  +  3(NH4)2S->2(NH4)3SbS4 
SnS2  +  (NH4)2S->(NH4)2SnS3 

This  difference  in  arsenic,  antimony,  and  tin  is  in  line  with 
the  general  observation  that  in  similar  groups,  metallic  character 
is  more  strongly  marked  as  atomic  weights  increase  (As=  75, 
Sb  =  120.2,  Sn=  119). 

Sub-group  A 

The  separation  of  the  sulphides  which  are  insoluble  in  yellow 
ammonium  sulphide  offers  a  number  of  interesting  considera- 
tions. Dilute  nitric  acid  dissolves  all  of  them  except  mercury 
sulphide,  because  by  reason  of  its  great  ionization,  it  reverses 
the  reaction,  MX  +  H2S  ^±  MS  +  HX  and  is  more  effective  than 
other  acids  because  it  removes  the  hydrogen  sulphide  as  formed 
and  so  facilitates  the  reversal : 

H2S  +  2  HNO,  ->  2  H9O  +  S  +  2  NOo 


QUALITATIVE  ILLUSTRATION  37 

In  case  the  nitric  acid  used  is  sufficiently  concentrated,  sulphur 
is  oxidized  and  the  sulphates  of  lead  and  mercury  may  be 
formed,  the  latter  being  soluble  and  the  former  fairly  insoluble 
in  the  acid  solution. 

The  sulphide  of  mercury  is  converted  to  the  chloride  by 
means  of  aqua  regia  which,  naturally,  must  be  completely  re- 
moved before  the  characteristic  test  for  mercuric  ion  is  made,- 
since  the  aqua  regia  also  converts  Sn++  to 

Hg++  +  Sn++  -+  Hg  + 
or  2  Hg++  +  Sn"+  ->  Sn++++  +  2  Hg++ 

The  filtrate  from  the  sulphide  of  mercury  is  treated  with  sul- 
phuric acid  to  convert  lead  nitrate  to  the  insoluble  sulphate,  and 
the  removal  of  nitric  acid  by  evaporation  is  essential,  if  all  the 
lead  is  to  be  so  converted,  since  the  reaction  is  reversible. 


Pb(N03)2  +  H2S04  -+  PbS04  +  2 

The  fact  that  no  lead  sulphate  precipitate  is  formed  until  the 
concentrated  acid  solution  is  diluted  is  due  to  the  primary  ion- 
ization  of  the  acid,  giving  rise  to  the  soluble  acid  salt,  which  in 
turn  is  decomposed  by  dilution. 

H2S04->H  +  HS04 
Pb++  +  2  HSO4  ^±  Pb(HSO4)2 
Pb(HS04)2  +  H20  ->  PbS04  +  4  H  +  2  S04 
The  filtrate  from  the  lead  sulphate  is  rendered  alkaline  by 
means  of  ammonium  hydroxide  and  the  bismuth  ion  so  con- 
verted into  bismuth  hydroxide.     The  mere  formation  of  a  floc- 
culent  precipitate  at  this  point  is  not  necessarily  indicative  of 
bismuth  because  of  the  presence  of  aluminium  salts  and  of  solu- 
ble silicates  in  ordinary  laboratory  reagents,  since  both  alumin- 
ium hydroxide  and  silicic  acid  are  precipitated  at  this  point. 

NOTE.    The  student  is  asked  to  explain  why  ammonium  silicate,  soluble  in  water,  is 
not  formed.     (  Vide  infra  '  Hydrolysis  '  and  Smith's  General  Chemistry,  p.  523.) 


38  QUALITATIVE  ANALYSIS 

The  confirmation  of  the  presence  of  bismuth  is  due  to  the 
hydrolysis  (vide  infra)  of  bismuth  chloride  and  the  insolubility 
of  the  bismuth  oxychloride. 

BiCl3  +  2  HOH  ±£  Bi(OH)2Cl  +  2  HC1 

This  reaction  is  easily  reversible,  and  the  presence  of  an  excess 
of  hydrochloric  acid  is  prevented  by  evaporation  or  by  addition 
of  sodium  acetate. 

NOTE.  The  student  is  asked  to  explain  how  the  latter  produces  this  result.  The  fol- 
lowing experiment  beautifully  illustrates  the  reversibility  of  the  reaction. 

EXPERIMENT  XXIV.  In  a  small  beaker  dissolve  some  Bids  in  the  small- 
est amount  of  concentrated  HC1  which  will  effect  solution.  Then  reverse  the 
reaction  by  addition  of  water  and  then  successively  by  HC1  and  water  as  long 
as  desired. 

The  ammonium  hydroxide  fails  to  precipitate  either  copper  or 
cadmium  hydroxides.  That  this  failure  to  precipitate  is  not  due 
to  the  amphoteric  nature  of  these  hydroxides  is  shown  by  the 
fact  that  they  are  not  dissolved  by  the  alkali  hydroxides.  The 

+  +  +  + 

complex  ions  Cu(NH3)4  and  Cd(NH3)4  form  hydroxides  solu- 
ble in  water,  which  are  in  equilibrium,  thus : 

Cd++  +  4  NH3  +  2  OH  ±£  Cd(NH3)4(OH)2 
Cu++  +  4  NH3  +  2  OH  ^±  Cu(NH3)4(OH)2. 

If  to  this  system  we  add  potassium  cyanide,  a  curiously  com- 
plex action  takes  place,  which  may  be  formulated  as  follows : 

Cd(+NH3)4^±4NH3 -h  Cd++ +  4  CN  ^±  Cd  (CN)4+2  K+ 
Cu"(NH3)4 ^± 4  NH3  +  Cu++  +  4  CN  ^±  Cu  (CN)4  +  2  K+. 

It  is  plain  that  the  dissociation  of  the  complex  negative  ions  is 
much  smaller  than  that  of  the  complex  positive  ions,  since  the 
point  of  rest  is  where  nearly  complete  transformation  in  the  di- 
rection (->•)  ensues,  as  is  shown  for  the  copper  ions  at  least  by 
the  removal  of  the  blue  color.  That  this  is  true  may  also  be 


QUALITATIVE  ILLUSTRATION  39 

shown  by  passing  hydrogen  sulphide  into  the  blue  solution, 
when  both  copper  and  cadmium  sulphides  are  precipitated.  If 
now  we  pass  hydrogen  sulphide  into  the  colorless  solution,  we 
find  a  concentration  of  Cd++  ions  sufficient  to  exceed  the  solu- 
bility product  for  cadmium  sulphide,  i.e.  CdxS>CdS.  The 
cadmium  sulphide,  instead  of  copper  sulphide,  is  precipitated  in 
spite  of  its  greater  solubility  product,  as  shown  in  Experiment 

XXI,  p.  33.     It  follows,  then,  that  Cd(CN)4  is  dissociated  to  a 
1 1 

greater  extent  than  Cu(CN)4. 

Sub-group  B 

The  reprecipitation  of  the  sulphides  of  arsenic,  antimony,  and 
tin  is  effected  by  dilute  acids  for  the  reason  that  the  free  acids 
H3AsS4,  H3SbS4,  and  H2SnS3  are  unstable  acids  of  the  charac- 
ter of  carbonic  acid,  and  break  down  into  hydrogen  sulphide 
and  As2S5,  Sb2S5,  and  SnS2.  The  sulphides  of  antimony  and 
arsenic  immediately  decompose,  forming  the  trisulphides.  Con- 
centrated hydrochloric  acid  dissolves  the  sulphides  of  tin  and 
antimony,  but  fails  to  dissolve  the  arsenic  sulphides.  The 
arsenic  sulphides,  when  oxidized  by  powerful  oxidizing  agents, 
yield  the  arsenic  oxide  which  forms  the  strong  acid  H3AsO4 
precipitated  by  magnesium  mixture  as  magnesium  ammonium 
arsenate. 

The  separation  of  antimony  and  tin  by  the  precipitation  with 
zinc  and  platinum  is  due  to  their  marked  difference  of  potential. 

NOTE.  The  alternative  method  of  detection  for  tin  and  antimony,  given  on  page  78,  is  to 
be  preferred  for  large  classes,  because  of  the  high  cost  of  platinum,  and  the  effort  is  made 
to  supply  a  satisfactory  method  of  procedure  which  will  make  the  use  of  platinum  appara- 
tus unnecessary. 

It  is  worthy  of  note  that  if  an  arsenate  is  present  in  the  orig- 
inal mixture  subjected  to  analysis,  the  negative  sulphur  ion  will 
not  react  with  the  negative  arsenate  ion,  and  in  order  that  pre^ 
cipitation  of  the  sulphide  may  take  place,  other  reactions  must 
precede.  These  reactions  may  be  expressed  as  follows  : 


40  Q  U  A  LIT  A  TIVE  ANAL  YSIS 

Na8AsO4+  3  HCl^±H8AsO4  +  3  NaCl 


As2O5  +  H2S  ^±  As2O3  +  H2O  +  S 
As2O3  4-  6  HC1  ^  2  AsCl3  +  3  H2O 
2  AsCl3  +  3  H2S  ^t  As2S3  +  6  HC1 

These  reactions,  with  the  exception  of  the  last,  being  easily 
reversible,  it  is  apparent  that  they  will  be  facilitated  by  in- 
creased concentration  of  hydrochloric  acid,  and  consequently 
strong  acidification  is  desirable,  if  arsenates  are  suspected. 
The  strong  acidity  of  the  solution  is  diminished  by  evaporation 
or  large  dilution  before  the  precipitation  is  completed,  for  the 
reasons  given  on  page  34. 

Group  III 

The  members  of  this  group  are  characterized  by  the  fact  that 
when  in  solution  in  such  forms  that  the  metals  furnish  positive 
ions,  they  are  all  precipitated  by  ammonium  sulphide  and  not 
by  H2S  in  acid  solution,  while  the  members  of  Groups  IV  and  V 
are  not  so  precipitated. 

These  elements  form  a  complicated  series  of  cations  and 
anions  not  all  of  which  are  likely  to  be  encountered  in  ordinary 
qualitative  operations,  but  which  may  be  tabulated  roughly  as 
follows  : 

CATIONS  ANIONS 

Co++     Co+++  Co(CN)6          Co(CN)6' 

Ni++     Ni+++  Ni(CN)6  Ni(CN)6 

Zn++  Zri62 


Mn++  Mn+++  Mn++++         Mn4    Mn4  MnO4 

Al+++  A1O3    A1O2 

Fe++  Fe+++  FeO3    FeO2    FeO4 

Cr++  Cr+++  Crb's    CrO2     CrO4  Cr2O7  Cr2O8(?) 


QUALITATIVE  ILLUSTRATION  41 

The  failure  of  the  positive  ions  to  precipitate  with  hydrogen 
sulphide  is  due  in  part  to  the  relatively  large  solubility  product 
of  the  sulphides  and  in  part  to  hydrolysis.  For  example,  the 
following  reaction  takes  place  as  indicated  though  the  precipita- 
tion is  not  complete : 

Zn(C2H302)2  +  H2S-*ZnS  +  2  HC2H3O2 

I 

The  reaction  may  be  driven  farther  toward  completion  by  the 
addition  of  sodium  acetate. 

'  NOTE.  The  student  is  asked  to  compare  this  effect  of  the  acetate  with  that  noted  on 
p.  35  (Experiment  XXIII).  Explain  and  formulate  the  explanation  by  a  series  of  equa- 
tions. (Cf.  p.  53,  on  the  solubility  of  Mg(OH)2  in  NH4C1.) 

In  the  presence  of  hydrochloric  acid  the  relatively  large  con- 
centration of  hydrogen  ions  prevents  any  precipitation  of  zinc 
sulphide  by  reason  of  the  suppression  of  the  sulphur  ions  of 
hydrogen  sulphide.  If  ammonium  sulphide  is  used,  we  have  a 
very  much  larger  concentration  of  sulphur  ions,  since  salts  are 
in  general  highly  ionized,  and  consequently  the  solubility  prod- 
uct of  zinc  sulphide  is  exceeded,  even  when  very  minute  quan- 
tities of  zinc  remain  in  solution,  for  since  Zn++  x  S  =  K,  when 

1 1 
the   concentration   of    S    is  small,   as   in   H2S    solutions,  then 

(Zn^+X^S)  =  K,  when  Zn  is  small,  if  X  is  very  large.  The 
same  considerations  apply  to  the  other  members  of  the  group 
with  modification  for  ions  of  chromium,  aluminium,  and  ferric 
iron.  Ferric  ions  as  well  as  the  negative  oxygen  complexes  of 
the  other  metals  do  not  exist  in  the  solution  for  analysis  of  this 
group,  if  previously  tested  for  the  second  group,  because  of  the 
reducing  properties  of  hydrogen  sulphide;  and  if  the  iron  has 
been  oxidized  by  means  of  nitric  acid,  it  is  again  reduced  by 
ammonium  sulphide  and  is  always  precipitated  as  the  sulphide, 

thus: 

2  FeCl3  +  (NH4)2S->2  FeCl2  +  2  NH4C1  +  S 

FeQ2  +  (NH^S-^FeS  +  2  NH4C1 


42  QUALITATIVE  ANALYSIS 

Chromium  and  aluminium  are  not  so  reduced,  yet  are  not  pre- 
cipitated as  trivalent  sulphides,  but  are  converted  to  hydroxides, 

thus: 

2  AlClg  +  3  (NH4)2S  ->  A12S3  -f  6  NH4C1 

-f  6  H2O  -»  3  H2S  +  2  A1(OH3) 


Hydrolysis 

The  hydrolytic  effect  of  water  illustrated  above  plays  a  most 
important  role  in  qualitative  analysis  and  has  already  been  en- 
countered in  the  behavior  of  bismuth  and  antimony  chlorides 
(cf.  reactions  52,  p.  67  and  118,  p.  75).  The  far-reaching  effect 
of  hydrolysis  is  well  illustrated  by  the  following  experiment, 
which  will  also  serve  as  a  basis  for  the  discussion  of  the  subject. 

EXPERIMENT  XXV.  Make  solutions  of  each  of  the  following  substances, 
NaCl,  Na2SO4,  CuSO4,  Na2CO3,  NaHCO3,  FeCl3.  Test  each  solution  with 
red  and  blue  litmus  paper. 

It  will  be  seen  that  the  behavior  of  'neutral'  salts  toward 
litmus  varies  widely  with  composition.  The  explanation  may 
be  had  from  a  consideration  of  the  law  of  mass  action  and  the 
ionization  '  constants  '  of  the  bases  and  acids  (cf.  p.  30).  The 
case  of  ferric  chloride  may  be  taken  as  a  type.  Since  it  is  a 
salt,  the  ionization  is  large.  Water  is  slightly  ionized,  therefore, 
in  addition  to  the  two  equilibrium  reactions, 


the  equilibrium  reactions  for  ferric  hydroxide  and  hydrochloric 
acid  must  also  be  produced. 


Since  the  ionization  constant  for  HC1  is  very  large  and  for 
Fe(OH)3  small,  then  the  original  balance  of  hydrogen  and  hy- 
droxyl  ions  is  disturbed,  and  an  excess  of  hydrogen  ions  results. 
It  follows  that  the  solution  is  acid  toward  litmus.  It  also  fol- 


QUALITATIVE  ILLUSTRATION  43 

i 
lows  that  since  OH  ions  are  withdrawn  from  solution  to  form 

un-ionized  Fe(OH)3,  that  water  must  continue  to  ionize. 


Eventually,  however,  in  spite  of  the  withdrawal  of  OH  ions 

we  have  '*  +  l)  \° l  ~y)  =  g^  ancj  the  water  ceases  to  ionize. 
//20 

This  result  is  apparently  reached  before  the  solubility  product 
Fe+++  x  (OH)3  =  K)  is  exceeded,  since  no  ferric  hydroxide 
precipitates. 

NOTE.  The  student  is  asked  to  explain,  in  a  similar  manner,  the  behavior  of  the  other 
solutions. 

In  case  of  salts  where  the  operation  of  hydrolysis  produces 
either  insoluble  or  slightly  ionized  substances  from  both  com- 
ponents, the  withdrawal  of  both  hydroxyl  and  hydrogen  ions 
permits  the  continuation  of  the  ionization  of  water  to  a  greater 
degree  and  frequently  even  to  complete  decomposition  of  the  salt. 

EXPERIMENT  XXVI.  —  To  a  fairly  concentrated  solution  of  aluminium 
sulphate,  add  a  strong  solution  of  sodium  carbonate,  and  note  the  formation 
of  the  precipitate  and  its  change  of  character  as  the  evolution  of  carbon  dioxide 
proceeds.  It  is  evident  visually  that  the  reaction  is  at  first  a  simple  metath- 
esis followed  by  the  hydrolysis  of  the  carbonate,  this  pair  of  reactions  being 
analogous  to  those  of  aluminium  sulphide,  on  p.  42. 

A  little  consideration  will  now  make  clear  to  the  student  that 
the  reaction  of  salts  in  solution  toward  litmus  is  a  function  of 
the  relative  strengths  (or  degree  of  ionization)  of  the  acids  and 
bases  from  which  they  may  be  considered  to  have  been  formed. 
We  may  then  distinguish  four  classes  of  salts  according  as  they 
are  formed  from  the  neutralization  of : 

1.  Strong  acid  and  strong  base  —  neutral  in  solution. 

2.  Strong  acid  and  weak  base  —  acid  in  solution. 

3.  Weak  acid  and  strong  base  —  alkaline  in  solution. 

4.  Weak  acid  and  weak  base  —  decomposed  in  solution. 


44  QUALITATIVE  ANALYSIS 

It  will  also  be  apparent  that  a  more  delicate  indicator  would 
serve  to  detect  a  difference  of  relative  strength  not  shown  by 
the  relatively  unsensitive  litmus. 

Since  aluminium,  chromium,  and  ferric  iron  are  all  very  weak 
base-forming  elements,  we  find  their  salts  with  the  weaker  acids 
are  all  hydrolized  by  water.  This  consideration  then  explains 
the  precipitation  of  the  hydroxides  of  the  elements  by  sulphides, 
carbonates,  and  acetates.  With  the  latter  salts  we  find  less 
rapid  hydrolysis,  as  would  be  expected,  considering  the  strength 
of  acetic  acid  as  compared  with  carbonic  acid  ;  and  in  the  case 
of  iron  acetate  the  hydrolysis  is  not  complete,  the  basic  acetate 
being  formed  probably  because  of  its  great  insolubility  rather 
than  because  of  the  stability  of  the  acetate.  These  considera- 
tions not  only  account  for  the  precipitation  of  the  group,  but 
also  for  the  separation  of  aluminium,  iron,  and  chromium  from 
manganese  and  zinc. 

Cobalt  and  Nickel  Reactions.  —  The  failure  of  cobalt  and 
nickel  sulphides  to  dissolve  in  dilute  hydrochloric  acid  is  remark- 
able in  view  of  the  fact  that  the  reaction 


CoCl2  +  H2S^±CoS  +  2  HC1 

reaches  equilibrium  without  precipitation  of  the  sulphide,  yet  if 
the  sulphide  is  formed,  then  a  much  greater  concentration  of 
acid  than  is  sufficient  to  prevent  its  formation  is  required  for 
solution.  The  explanation  is  probably  to  be  found  in  the  poly- 
merization of  the  sulphides  after  precipitation  forming  (CoS)^ 
and  (NiS)x  which  are  more  insoluble  than  the  simple  molecules. 
The  doubt  in  this  case  might  be  resolved,  had  we  a  method  of 
determining  the  molecular  weight  of  such  solids. 

The  remarkable  similarity  of  properties  of  cobalt  and  nickel, 
which  is  shown,  not  alone  in  the  metallic  state,  but  also  in  their 
ionic  reactions,  is  responsible  for  the  somewhat  involved  reactions 
necessary  for  their  separation.  The  peculiar  relation  of  these 


QUALITATIVE  ILLUSTRATION  45 

elements,  the  atomic  weights  of  which  are  so  nearly  equal,  Ni 
58.68,  Co  58.97,  is  nowhere  more  strikingly  shown  than  in  the 
following  experiment : 

EXPERIMENT  XXVII.  Make  solutions  of  cobalt  chloride  and  nickel 
chloride  and  place  a  portion  of  each  to  one  side  for  later  comparison.  The 
bulk  of  each  solution  is  treated  with  a  small  excess  of  a  saturated  solution  of 
sodium  bicarbonate  and  the  precipitate  allowed  to  settle.  The  supernatant 
liquid  is  poured  off,  and  both  precipitates  are  cooled  with  ice  and  salt  and 
treated  with  a  small  excess  of  hydrogen  peroxide,  which  is  also  cooled  in  ice 
and  salt.  The  two  solutions  are  now  filtered  quickly  and  compared  in  color 
with  the  original  pink  cobalt  and  green  nickel  solutions.  The  colors  will  be 
found  to  be  interchanged,  the  cobalt  solution  is  green,  the  nickel  pink. 

NOTE.  The  nickel  solution  should  be  filtered  through  an  ice-cold  filter,  but  even  so 
the  pink  color  is  very  evanescent,  the  green  color  returning  with  simultaneous  evolution 
of  oxygen. 

Because  of  the  close  similarity  of  the  ordinary  salts  of  these 
elements,  recourse  is  had  to  the  slightly  greater  differences 
shown  by  them  when  in  the  trivalent  form  and  as  portions  of 
complex  negative  ions.  The  series  of  reactions  is  identical  for 
both  elements,  thus  : 

CoCL,  +  2  KCN  ->  Co(CN)2  +  2  KC1 
Co(CN)2  +  4KCN  ->  K4Co(CN)6 
K4Co(CN)6  +  Br->KBr  +  K3Co(CN)6 

There  is,  however,  a  distinction  between  the  double  cobalti- 
cyanide  and  the  nickel  cyanide.  Both  are  in  equilibrium  sys- 
tems as  follows : 

Co+^  +  3  CN  ^±  Co(CN)3  +  3  KCN  ^±  K3Co(CN)6 

^±3K+  +  Co'(CN)6 
3  CN  ^±  Ni(CN)3  +  3  KCN  ^±  K3Ni(CN)6 


46  QUALITATIVE  ANALYSIS 

If  to  the  systems  an  excess  of  sodium  hydroxide  is  now  added, 
the  product  of  the  concentration  of  the  Ni+++  ions  and  the  OH 
ions  exceeds  the  solubility  product  for  Ni(OH)3,  but  not  for 
Co(OH)3.  The  inference  is,  of  course,  that  the  double  cobalt 
salt  is  the  more  stable  compound.  A  similar  slight  difference 
of  the  double  nitrite  ions  is  found  sufficient  for  separation  since 
K3Co(NO2)6  is  insoluble  in  acetic  acid,  while  the  K3Ni(NO2)6  is 
soluble. 

NOTE.  For  the  methods  of  distinguishing  cobalt  and  nickel  by  means  of  the  reaction 
of  Tschugaeff,  see  Ber.  d.  Chern.,  38  :  2520. 

Manganese  and  Zinc  Separation.  The  separation  of  zinc 
from  manganese  depends  upon  the  greater  solubility  product  of 
the  manganese  sulphide,  so  that  while  in  acetic  acid  solution  it 
is  easily  possible  to  obtain 

Zn  x  S  >  KK,  ZnS 
but 

Mn  x  S  <  K^m  MnS. 


Manganese  is  by  powerful  oxidation  converted  to  the  highly 
colored  MnO4  ion.  The  principles  of  oxidation  will  be  dis- 
cussed in  connection  with  chromium. 

Separation  of  Iron,  Aluminium,  and  Chromium.  As  was  ob- 
served on  p.  35  in  connection  with  the  arsenic  group,  increase 
of  valency  is  in  general  accompanied  by  increased  tendency  to 
form  negative  complex  ions.  This  has  shown  itself  in  the  for- 
mation of  the  negative  complexes  of  trivalent  cobalt  and  nickel 
and  also  in  the  negative  character  of  the  MnO4  ion,  where  the 
manganese  is  heptavalent  as  compared  with  the  strongly 
metallic  character  of  Mn++.  Also  on  p.  36  attention  was 
directed  to  decreasing  acidic  tendencies  in  similar  groups  with 
increasing  atomic  weights  (Al  =  2/;  0=52;  Fe=56). 
Both  these  characteristics  are  manifested  in  the  group  under 
consideration. 


QUALITATIVE  ILLUSTRATION  47 

Aluminium  hydroxide  is  a  very  striking  example  of  an  am- 
photeric  substance,  i.e.  one  which  is  at  the  same  time  both  acidic 
and  basic.  This  is  shown  by  the  following  reactions. 

A1(OH8)  4-  3  HCl-»  AlClg  +  3  H2Q 
A1(OH)3  +  3  NaOH  -+  Na3AlO3  +  3  H2O 

Written  ionically  these  become 

A1+++  +  3C1->A1C13 
A163  +  3Na->Na3AlO3. 

Aluminium  hydroxide  in  solution  must  therefore  satisfy  two  con- 
stants, viz.  : 


al(ok\ 

where  both  K  and  K'  are  extremely  small,  i.e.  it  is  both  a  weak 
acid  and  weak  base,  and  by  consequence  any  tendency  to  form 
salt  by  autoneutralization  is  offset  by  hydrolysis.  Moreover, 
the  solubilit^  product  is  also  small  so  that  but  minute  quantities 
are  in  solution.  The  system  may  perhaps  best  be  illustrated  by 
the  following  equilibrium  reaction  : 

3  H+  4-  Al'63±^  A1(OH)3^±A1+++  +  3  OH 

It 
A1(OH)3  (solid) 

If  now  to  this  system  we  add  hydroxyl  ions  by  means  of  a 
strong  -  base,  there  results  the  withdrawal  of  hydrogen  ions  to 
form  water  and  the  highly  ionized  aluminate,  and  the  whole 
system  moves  in  the  direction  indicated  by  (<—).  If,  on  the 
other  hand,  we  add  a  strong  acid,  hydroxyl  ions  are  removed 
and  the  system  moves  in  the  direction  (—>•).  Ammonium  hy- 
droxide is  too  weakly  basic  to  affect  the  system  materially,  and 
they  same  statement  is  true  for  the  very  weak  acids.  The  fact 
that  acetic  acid  (approximately  ionized  to  the  same  extent  as 
ammonium  hydroxide)  does  dissolve  the  hydroxide  indicates 


48  QUALITATIVE  ANALYSIS 

that  more  hydroxyl  ions  are  present  than  hydrogen  ions,   i.e. 
A1(OH)3  is  more  strongly  basic  than  acidic. 

Chromium  hydroxide  acts  in  a  similar  manner,  but  is  more 
weakly  acid  than  aluminium  hydroxide,  and  therefore,  while  it 
does  dissolve  to  some  extent  in  sodium  hydroxide  in  the  cold, 
its  salt  is  completely  hydrolyzed  by  boiling,  and  the  hydroxide 
is  reprecipitated.  Ferric  hydroxide  shows  so  little  acid  char- 
acter as  to  be  almost  wholly  unaffected  by  alkali  solvents. 
Hence,  the  separation  of  iron  and  chromium  from  aluminium 
hydroxide  and  the  identification  of  the  latter. 

NOTE.  Both  iron  and  chromium  hydroxides  on  fusion  with  bases  form  salts  of  the 
corresponding  meta  acids  even  as  aluminium  hydroxide  does.  This  class  of  compounds 
is  abundant  in  nature  and  is  known  as  the  spinels,  examples  of  which  are 

Fe(CrO2)2    chromite          Zn(FeO2)2    franklinite 
Mg(AlO)2     spinel  Fe(FeO2)2    magnetite 

Iron  and  chromium  both  form  distinct  acids  in  a  higher  state 
of  oxidation.  Here  also  the  difference  of  degree  manifests 
itself  clearly.  Fusion  with  sodium  carbonate  and  potassium 
nitrate  converts  the  chromium  hydroxide  completely  to  a  chro- 
mate,  and  the  iron  hydroxide  remains  unaffected  except  by  con- 
version to  the  oxide  Fe2O3  by  dehydration.  In  the  presence  of 
hydrogen  peroxide,  the  chromium  hydroxide  may  be  converted 
to  the  chromate  at  ordinary  temperature  in  solution.  However, 
fusion  with  sodium  peroxide  converts  the  iron  hydroxide  into 
unstable  ferrates,  which  are  soluble  in  water  but  decompose  on 
boiling.  In  slightly  acid  solutions  chromates  may  be  oxidized 
to  a  still  greater  extent,  forming  the  free,  but  very  unstable, 
acid  H2Cr2O8,  the  anhydride  of  which,  Cr2O7,  is  soluble,  and 
somewhat  more  stable,  in  ether. 

Oxidation 

The  r61e  played  by  oxidation  in  the  separation  of  the  metals 
is  very  important,  and  a  brief  discussion  of  the  general  subject 
is  in  order.  By  derivation  the  meaning  of  the  term  is  limited  to 


QUALITATIVE  ILLUSTRATION  49 

such  actions  as  involve  the  union  of  oxygen  with  some  other 
substance.  Examples  are,  of  course,  familiar.  In  the  character 
of  these  changes  there  is,  however,  no  essential  difference  from 
those  in  which  union  with  elements  other  than  oxygen  is  effected. 
The  use  of  the  term  has  thus  become  extended  to  cover  those 
cases  where  an  element  changes  its  valency  in  a  positive  direc- 
tion. Thus  iron  is  oxidized  by  chlorine  to  form  ferrous  chloride, 
and  this  in  turn  is  further  oxidized  to  ferric  chloride,  though  no 
oxygen  is  involved  in  the  operation.  Considered  from  the  ionic 
viewpoint  any  change  in  the  electrical  charge  upon  the  ions 
which  increases  the  positive  charge  or  diminishes  the  negative 
charge  is  an  oxidation.  For  example,  the  ferrocyanide  ion  is 
oxidized  to  the  ferricyanide  by  loss  of  one  negative  charge  (or 
electron).  (See  Walker's  Physical  Chemistry,  p.  307.)  While 
it  is  helpful  to  look  upon  change  of  ionic  charge  as  the  essential 
factor  in  oxidation  and  reduction  of  ions,  and  while  it  is  possible 
also  to  interpret  the  formation  and  decomposition  of  complex 
ions  in  these  terms,  it  is  somewhat  simpler  to  view  these,  at 
times  somewhat  complex,  changes  only  from  the  standpoint  of 
change  in  valency. 

In  the  fusion  reaction  using  potassium  nitrate,  there  is  direct 
addition  of  oxygen  in  the  case  of  both  manganese  and  chro- 
mium. It  is  not  possible  to  obtain  the  products  of  oxidation  in 
either  case  in  the  absence  of  a  basic  material,  for  the  reason 
that  the  free  oxides  are  unstable  at  the  temperatures  required 
for  their  formation.  We  may  formulate  the  changes  as  follows  : 


2  Cr(OH)3  ->  H2O+  Cr2O3  +  3  O  ^  2  CrO 
CrO3  +  Na^Og  ^  Na2CrO4  +  CO2 


The  sodium  carbonate  also  plays  the  part  of  a  flux,  or  liquid 
solvent,  rendering  contact  of  the  interacting  materials  more  inti- 


50  QUALITATIVE  ANALYSIS 

mate  and  thus  facilitating  the  reaction.  All  these  changes  may 
be  thrown  into  ionic  form  but  are  not  illuminated  greatly  by  the 
operation. 

Bead  Tests 

Some  of  the  metal  ions  of  the  third  as  well  as  of  the  second 
group  are  capable  of  detection  by  the  so-called  bead  reactions, 
the  indications  of  which  are  tabulated  on  p.  124.  The  bead  reac- 
tions are  essentially  dependent  upon  the  conversion  of  the  salts 
of  meta  acids  into  corresponding  ortho  salts.  The  simplest 
example  is  furnished  by  the  microcosmic  salt  reactions,  as  illus- 
trated by  the  following  equations  : 

NaNH4HPO4  ->  NaH2PO4 
NaH2PO4->NaPO3  +  I 

If  we  add  to  the  bead,  NaPO3,  a  non-volatile  oxide,  in  small 
quantity,  the  oxide  will  be  practically  wholly  dissolved,  thus : 

NaPO3  +  M2O  -^  NaM2PO4 

If  the  metallic  oxide  is  bivalent  or  trivalent,  the  character  of 
the  reaction  is  unchanged,  though  the  formulas  of  the  com- 
pounds may  be  more  complex.  If  borax  is  used  to  form  the 
bead,  there  is  no  essential  variation,  and  the  borax  glass  may 
be  considered  as  a  mixture  of  the  metaborate  and  the  boric 
anhydride. 

Na2B4O7  -  10  H2O  -+  Na2B4O7  +  10  H2cf 
Na2B4O7->2  NaBO2  4-  B2O3 
NaBO2  +  CuO  ->  NaCuBO3 

All  metallic  oxides  undergo  these  reactions,  but  not  all  furnish 
characteristically  colored  salts.  The  different  valences  of  the 
metals  also  vary  in  the  color  of  salt  formed.  Non-metallic 
oxides  are  of  course  incapable  of  playing  the  part  indicated 


QUALITATIVE  ILLUSTRATION  51 

above,  but  with  a  basic  bead,  as  Na^Og,  they  react  in  an  anal- 
ogous manner,  thus  : 

Na2CO3  +  SiO2->  Na2SiO3  4-  CO2. 


Group  IV 

The  members  of  Group  IV  are  characterized  by  being  in- 
capable of  precipitation  from  solution  by  HC1,  H2S,  or  (NH^S 
and  by  formation  of  carbonates  insoluble  in  ammonium  salts. 
Only  bivalent  cations  are  encountered.  The  solution  of  the 
group  precipitate  is  best  effected  by  weak  acetic  acid,  so  that, 
on  treatment  with  chromate,  the  reversal  of  the  reaction  is  less 
vigorous  than  with  a  stronger  acid. 

2  Ba(C2H3O2)2  +  K2Cr2O7  4-  H2O  ->  2  BaCrO4  +  2  KC2H3O2 

\        +  2  HC2H302 

After  removal  of  the  surplus  dichromate  by  reprecipitation 
and  resolution  in  acetic  acid,  the  test  for  the  presence  of  stron- 
tium ions  is  made  on  a  portion  of  the  solution  by  addition  of  a 
saturated  solution  of  calcium  sulphate.  Since  calcium  sulphate 
is  soluble  only  to  the  extent  of  2  grams  per  liter  (see  Table  of 
Solubilities,  p.  174),  it  is  evident  that  the  precipitate  of  strontium 
sulphate  can  only  be  very  slight,  regardless  of  the  amount  of 
strontium  present.  This  precipitate  may  not  appear  at  once 
because  of  supersaturation  (q.v.,  p.  55).  The  portion  not  used 
for  the  strontium  test  is  treated  with  sodium  sulphate  to  remove 
the  strontium  ions,  and  since  calcium  sulphate  is  so  sparingly 
soluble,  all  but  two  grams  per  liter  will  precipitate  with  the 
strontium.  Consequently,  on  treatment  with  ammonium  oxalate, 
only  a  very  slight  precipitate  of  calcium  oxalate  is  to  be  expected. 

The  most  interesting  feature  of  this  group,  which  also  fur- 
nishes an  interesting  study  of  relative  solubilities,  is  presented 
by  the  filtrate  from  Group  IV.  From  the  Table  of  Solubilities, 
it  will  be  seen  that  the  following  quantities  of  carbonates  remain 


52  QUALITATIVE  ANALYSIS 

in  solution,  if  precipitation  of  each  has  taken  place :  SrCO3, 
ii  mg.  per  liter;  CaCO3,  13  mg.  per  liter;  BaCO3,  23  mg.  per 
liter.  Small  as  are  these  quantities,  a  very  satisfactory  separa- 
tion and  identification  is  possible  in  this  filtrate.  Treatment 
with  dilute  sulphuric  acid  should  give  a  precipitate  of  barium 
sulphate  (solubility  2.3  mg.  per  liter),  and  when  this  is  filtered, 
the  filtrate  with  ammonium  hydroxide  and  ammonium  oxalate 
should  give  a  precipitate  of  calcium  oxalate  (solubility  5.6  mg. 
per  liter).  No  reagent  will  give  a  precipitate  of  strontium  of  a 
solubility  corresponding  to  less  than  1 1  mg.  per  liter,  but  evap- 
oration to  dryness  and  expulsion  of  the  ammonium  salts  will 
furnish  a  residue  suitable  for  a  flame  test. 

Flame  Tests 

Since  each  of  the  ions  of  this  group  furnishes  a  characteristic 
flame  coloration,  it  should  be  noted  here  that  only  salts  which 
are  volatile  at  the  temperature  of  the  Bunsen  flame  can  be  used 
for  this  purpose.  It  is  therefore  necessary  to  convert  sulphates, 
chromates,  etc.,  into  the  more  volatile  chlorides  by  moistening 
with  hydrochloric  acid.  The  volatile  chlorides  then  dissociate 
in  the  flame,  and  the  ions  produce  color  waves  which  are  the  re- 
sultants of  their  spectral  lines.  Strontium  chloride  gives  a 
somewhat  evanescent  flame  because  of  its  rapid  volatilization. 

Group  V 

The  members  of  Group  V  form  what  is  called  the  soluble 
group  because  relatively  few  of  their  salts  are  insoluble  in  water. 
With  the  exception  of  magnesium,  they  form  univalent  cations 
in  aqueous  solution. 

Ammonium.  Solutions  of  ammonium  hydroxide  furnish  a 
variety  of  reactions  which  depend  not  only  upon  the  formation 
of  the  NH4  ion  by  electrolytic  dissociation,  but  also  upon  the 
dissociation 

H20  (i) 


QUALITATIVE  ILLUSTRATION  53 

This  latter  dissociation  manifests  itself,  at  somewhat  elevated 
temperatures,  with  ammonium  salts. 


:2  NH3  +  H2S04 

It  is  in  consequence  of  the  volatility  of  one  or  both  of  the  prod- 
ucts of  this  dissociation  that  the  most  convenient  test  for  am- 
monium is  to  heat  the  suspected  substance  with  a  base  and  when 

X 

the  reaction  NH4X  +  MOH  ->  MX  +  NH4OH  ->  NH3  +  H2O 

has  taken  place  to  reverse  reaction  (i)  by  absorption  of  the 
volatile  ammonia  in  moist  litmus  paper.  The  most  delicate  test 
for  the  ammonium  ion  does  not  depend  directly  upon  its  ionic 
nature.  Nessler's  reagent  reacts  with  ammonium  by  substitution 
of  mercury  for  hydrogen.  The  reaction  may  be  formulated  : 

NH4OH  +  2  K2HgI4  +  3  KOH  ->  NHg2I  +  7  KI  +  4  H2O 

Sodium.  Sodium  ions  form  no  very  insoluble  salts  and  there- 
fore dependence  is  usually  placed  upon  the  somewhat  embar- 
rassingly delicate  flame  test.  Attention  is  directed,  however,  to 
reactions  260  and  261  on  p.  96,  particularly  the  latter. 

Potassium.  Potassium  ions  form  a  few  insoluble  salts,  the  one 
most  frequently  used  for  identification  being  the  chlorplatinate. 
From  the  standpoint  of  economy,  the  precipitate  formed  by 
sodium  cobaltinitrite  is  better,  without  material  difference  in 
efficiency. 

Magnesium  :  The  presence  of  magnesium  in  this  group  is 
somewhat  anomalous  when  one  considers  its  place  in  the  periodic 
system  and  the  solubility  of  the  carbonate  (i  g.  per  liter)  and 
hydroxide  (lomg.  per  liter).  It  escapes  precipitation  in  both 
the  third  and  fourth  groups  by  reason  of  the  presence  of  ammo- 
nium salts.  The  reaction  between  magnesium  salts  and  ammo- 
nium hydroxide  is  easily  reversible  and  the  following  experiment 
is  very  instructive. 


54  QUALITATIVE  ANALYSIS 

EXPERIMENT  XXIX.  Make  a  concentrated  solution  of  MgCl2  and  treat 
with  concentrated  ammonium  hydroxide  and  filter.  To  the  precipitate,  add  a 
solution  of  ammonium  chloride  and  shake.  The  hydroxide  should  dissolve. 
To  the  filtrate,  add  sodium  phosphate  ;  the  magnesium  not  precipitated  by  the 
hydroxide  is  shown  to  be  considerable.  To  some  of  the  original  solution  add 
ammonium  chloride  and  then  ammonium  hydroxide  ;  no  precipitation  follows. 

This  rather  remarkable  reversion  is  usually  explained  as  fol- 
lows :  The  precipitation  of  Mg(OH)2  depends  upon  the  prod- 
uct of  the  ions  being  in  excess  of  the  solubility  product,  i.e. 
mg  x  olfi  >  A"Mg(0H)2-  Ammonium  hydroxide  is  a  weak  base 


•vide  infra,  and  being  in  equilibrium,    —  ^  —  —  =  K,  the  excess 

;z^4oh 

of  ammonium  ions  from  the  highly  dissociated  chloride  depresses 
the  ionization  to  the  practical  elimination  of  hydroxyl  ions.  The 
equilibrium  reaction  (somewhat  awkwardly  expressed) 


Mg++  +  2d  NH4  +  OH  Mg2(OH)  NH4  +  C1 

lit  +  H  5:          H  +  It 

MgCl2  2NH4OH  Mg(OH)2  2  NH4C1 

A| 

Mg(OH)2 
is  therefore  driven  in  the  direction  f  •<— t )  by  the  addition  of  am- 


-{)  by  tl 


monium  ions  and  their  presence  before  treatment  with  ammo- 
nium hydroxide  prevents  the  movement  in  the  direction  (— >)  to 
a  sufficient  extent  to  form  Mg(OH)2  in  excess  of  its  solubility. 
A  similar  explanation  accounts  for  the  failure  to  precipitate 
Zn(OH)2  and  Mn(OH)2  when  ammonium  hydroxide  and  am- 
monium chloride  are  added  to  the  third  group  solution  in  prep- 
aration for  treatment  with  ammonium  sulphide.  (For  failure  to 
precipitate  cobalt  and  nickel  salts,  see  p.  32.  Consult  also 
rules  of  solubility,  p.  101.) 

The  fact  that  ammonium  carbonate  does  not  precipitate  mag- 
nesium salts  in  the  presence  of  ammonium  salts  might  be  ascribed 


QUALITATIVE  ILLUSTRATION  55 

to  a  similar  elimination  of  CO3  ions  by  suppression,  were  it  not 
that  an  excess  of  a  common  ion  exercises  a  relatively  less  marked 
effect  on  a  highly  ionized  substance  than  upon  one  slightly  ion- 
ized. It  were  scarcely  possible  to  wholly  suppress  the  ionization 
of  ammonium  carbonate  by  means  of  ammonium  chloride.  It  is 
perhaps  preferable  to  consider  that  magnesium  ions  form  with 
the  chloride  the  double  salt  NH4MgCl3  (analogous  to  carnellite 
KMgClg),  which  is  not  decomposed  in  the  presence  of  the  CO3 
ion. 

The  ammonium  chloride  also  plays  a  role  in  the  precipitation 
of  magnesium  by  soluble  phosphates  in  ammoniacal  solution. 
The  reaction  may  be  expressed  thus : 

[MgCl2+  NH4OH  +  Na2HP04-*MgNH4P04  +  2NaCl+  H2O. 

.  * 

The  presence  of  the  ammonium  chloride  prevents  precipitation 
of  any  magnesium  hydroxide  by  weakening  the  basic  reaction 
of  the  hydroxide.  The  slow  formation  of  the  phosphate  precipi- 
tate is  due  to  supersaturation. 

Supersaturation 

Supersaturation  of  solutions  is  frequently  encountered  when 
the  substance  is  formed  in  solution,  or  when  substances  more 
soluble  in  warm  than  in  cold  solvents  are  dissolved  to  saturation 
at  a  given  temperature  and  then  brought  to  a  lower  tem- 
perature. The  following  experiment  beautifully  illustrates  the 
phenomenon  and  also  the  property  of  independent  solubility 
.  and  the  effect  of  the  presence  of  the  solid  substance. 

EXPERIMENT  XXX.  Partly  fill  a  test  tube  with  crystals  of  sodium  thiosul- 
phate  and  heat  until  dissolved  completely  in  the  water  of  crystallization.  Pour 
on  top  of  this  a  layer  of  a  solution  of  sodium  acetate  saturated  at  50°  C.  and 
allow  to  cool.  A  crystal  of  sodium  thiosulphate  may  now  be  dropped  through 
i  the  acetate  into  the  lower  layer  which  will  immediately  crystallize.  A  crystal 
of  sodium  acetate  may  now  be  dropped  into  the  upper  layer,  which  will  im- 


56  QUALITATIVE  ANALYSIS 

mediately  crystallize.  There  will  remain  a  middle  portion  which  is  not 
solidified  by  reason  of  the  mutual  dilution  of  the  solutions  by  each  other. 

The  experiment  illustrates  the  most  effective  means  of  caus- 
ing precipitation  in  supersaturated  solutions.     It   is    apparent 

^ 
that  from  the  equilibrium  formula,  —  =  K,  if  S'  =  o,  then  5  may 

ijj 

be  of  any  magnitude,  so  far  as  the  law  is  concerned.  If,  how- 
ever, a  portion  of  solid  substance  is  introduced,  the  material  in 
solution  must  alter  in  quantity  until  the  constant  ratio  is  reached. 
The  same  result  is  obtained  by  violent  agitation  or  by  scratching 
the  inner  glass  surface  with  a  rod.  The  principle  involved  is 
that  violent  agitation  initiates  the  formation  of  a  minute  quantity 
of  solid,  after  which  the  operation  continues  according  to  the  law 
of  physical  equilibrium.  Ordinarily  standing  for  a  considerable 
time  also  produces  the  discharge  of  supersaturation. 


PART    II 

METAL   ANALYSIS 
General  Directions. 

EQUATIONS.  Every  chemical  reaction  is  capable  of  more 
or  less  accurate  expression  in  the  form  of  an  equation.  This 
consists  of  a  representation,  by  means  of  atomic  symbols,  of  the 
component  parts  of  a  reaction  and  of  the  nature  and  direction  of 
the  change.  An  equation  therefore  indicates  the  quantitative 
relations  of  the  substances  in  so  far  as  reaction  occurs.  An 
equation  never  conveys  any  indication  of  the  condition  necessary 
to  produce  the  indicated  change  nor  the  completeness  of  its 
accomplishment.  Neither  does  an  equation  indicate  whether  or 
not  other  reactions  may  be  of  simultaneous  occurrence,  nor 
whether  the  change  may  be  immediately  followed  by  subsequent 
change  of  the  products  of  the  reaction.  Such  an  expression  is 
therefore  far  from  a  complete  exposition  of  a  chemical  change. 
Yet  equations  set  before  us  in  such  concise  form  the  character  of 
the  relations  with  which  we  deal  that  facility  in  their  use  is 
indispensable  in  a  study  of  qualitative  analysis.  The  student  is 
expected,  therefore,  to  write  equations  representing  the  chemical 
changes  involved  in  the  preliminary  reactions.  A  study  in  equa- 
tion making  is  found  in  Walker's  Physical  Chemistry^  pp.  22-27. 

PRECIPITATION  AND  FILTRATION.  The  part 
played  by  precipitation  in  qualitative  analysis  has  been  dis- 
cussed in  the  introduction.  The  importance  of  complete  pre- 
cipitation is  mentioned  on  p.  34.  It  remains  to  point  out  the 

57 


58  QUALITATIVE  ANALYSIS 

method  of  preparing  precipitated  substances  for  examination. 
The  precipitant  is  to  be  added  slowly,  the  substances  thrown  out 
of  solution  allowed  to  settle,  where  possible,  and  more  precipitant 
added  until  no  further  change  takes  place.  After  filtering,  a 
small  amount  of  precipitant  should  always  be  added  to  the  filtrate 
as  a  precautionajg^measure.  Wherever  possible  the  filtration 
should  be  from  hoi  solutions,  both  because  the  precipitates 
coagulate  more  readily  when  hot  and  settle  out  more  completely, 
and  because  hot  liquids  pass  through  filters  more  rapidly  than 
cold.  The  filtering  medium  is  usually  paper  of  a  bibulous  charac- 
ter, which  is  cut  into  circular  form  and  folded  to  form  a  quadrant. 
It  is  opened  between  the  third  and  fourth  layers  and  fitted  into 
a  funnel  with  the  help  of  water  from  a  wash  bottle.  If  the  paper 
be  fitted  by  its  upper  edge  into  the  funnel,  but  not  closely  along 
the  sides,  the  rate  of  filtration  is  materially  increased.  The  size 
of  the  funnel  should  be  such  that  the  paper  covers  about  two 
thirds  of  the  side.  The  liquid  to  be  filtered  is  poured  into  the 
paper  along  a  glass  rod  held  vertically,  and  the  pouring  should 
be  at  such  a  rate  that  the  paper  is  kept  continually  about  half  full. 
A  bit  of  soft  rubber  tubing  slipped  over  the  end  of  the  glass  rod 
furnishes  a  convenient  instrument,  a  'policeman,'  for  transferring 
the  last  portions  of  the  precipitate  into  the  filter. 

WASHING.  After  all  the  precipitate  is  thus  transferred  to 
the  paper,  it  must  be  washed  in  order  to  free  it  from  adhering 
mother  liquor  and  to  bring  into  the  filtrate  all  material  which  prop- 
erly belongs  in  it.  Wherever  possible,  washing  should  be  done 
with  hot  water,  for  which  a  wash  bottle  with  the  neck  wrapped  in 
cork  or  asbestos  is  a  great  convenience.  In  all  cases  washing 
should  be  accomplished  by  means  of  small  amounts  of  water 
added  repeatedly  rather  than  by  immediate  application  of  a 
large  quantity.  A  very  interesting  discussion  of  this  matter 
is  found  in  Ostwald's  Foundations  of  Analytical  Chemistry, 
pp.  1 1-22. 


METAL  ANALYSIS  59 

DECANT  ATION.  Where  precipitates  are  either  gelatinous 
or  heavy,  time  is  frequently  saved  by  allowing  the  predicate  to 
settle  and  transferring  the  supernatant  liquid  to  the  filter  without 
disturbing  the  precipitate.  The  latter  may  then  be  washed  in 
the  vessel  in  which  it  was  formed  and  th^jrah  water  again  de- 
canted through  the  filter.  In  this  manner  clogging  of  the  filter 
and  consequent  slow  filtration  may  be  avoided,  and  at  times  it 
may  not  be  necessary  to  transfer  the  precipitate  to  the  paper. 

EVAPORATION  OF  FILTRATES.  Where  filtrates  are 
to  be  concentrated  or  evaporated  to  dryness,  it  is  not  usually 
advisable  to  use  a  free  flame,  owing  to  the  danger  of  loss  by 
ebullition,  spattering,  or  decomposition  when  dry.  The  most 
satisfactory  apparatus  is  a  laboratory  water  bath,  but  where  this 
is  not  convenient,  a  lip  beaker,  preferably  of  porcelain,  may  be 
advantageously  substituted. 

AMOUNT  OF  SAMPLE.  The  beginner  in  qualitative 
analysis  finds  it  difficult  to  persuade  himself  that  greater  certainty 
is  not  reached  by  using  large  quantities  of  material.  Experience 
teaches  that  small  quantities  are  preferable,  not  alone  by  reason 
of  greater  ease  in  securing  complete  separations,  but  because 
much  time  is  saved  in  manipulation.  A  fair  general  rule  for  un- 
known samples  for  complete  analysis  is  to  divide  into  four  parts  ; 
one  for  preliminary  examination  (see  part  IV),  one  for  metal,  and 
one  for  acid  analysis,  while  the  fourth  is  a  reserve  which  in  case 
of  unsatisfactory  results  may  be  treated  as  the  original  sample. 
A  gram  of  substance  ought  ordinarily  to  be  sufficient  for  all  of 
these  analyses. 

CONFIRMATION  OF  TESTS.  In  the  final  identification, 
dependence  is  seldom  to  be  placed  upon  one  test  alone.  Each 
final  reaction  should  be  supplemented  by  additional  evidence, 
when  possible,  and  a  sharp  outlook  for  indications  should  be  kept 
at  all  stages  of  analysis. 

TESTS   OF   REAGENTS.     It  is   not   always   possible   to 


60  QUALITATIVE  ANALYSIS 

secure  perfectly  pure  reagents,  and  when  analysis  reveals  only 
traces  «f  a  particular  substance,  it  is  frequently  advisable  to  test 
the  reagents  used  for  this  substance.  This  is  particularly  true 
for  such  common  substances  as  iron,  aluminium,  sodium,  and 
chlorine.  This  examination  of  the  reagents  is  made  by  subject- 
ing the  reagents  used  to  the  same  series  of  tests  as  the  sample 
which  is  undergoing  examination.  This  is  called  "running  a 
blank."  In  case  the  substance  in  question  is  found  in  the  re- 
agents, judgment  must  be  passed  on  the  relative  amounts  found, 
in  case  purer  reagents  are  not  to  be  had. 

NOTEBOOKS.  Notes  are  to  be  kept,  in  which  are  to  be 
recorded,  besides  the  equations  already  mentioned,  the  results  of 
the  analyses  carried  out  and  the  answers  to..J:he  exercises  which 
are  found  after  each  group.  The  student  should  also  record  any 
special  observations  made  and  any  valuable  hints,  precautions,  or 
interesting  facts  which  are  picked  up  in  the  course  of  study  or 
work,  and  which  are  not  to  be  found  in  the  manual. 

Group  I 

The  Hydrochloric  Acid  Group 
Silver,  Mercury  (pus),  and  Lead 

GENERAL  STATEMENT.  The  metals  of  this  group  are 
precipitated  by  hydrochloric  acid  —  lead  incompletely.  If  lead 
is  found  in  this  group,  it  will  also  be  found  in  Group  IL 

SILVER  :  Solution  for  Reactions,  Silver  Nitrate. 

i.  Hydrochloric  acid  or  a  soluble  chloride  precipitates 
white  silver  chloride,  AgCl ;  soluble  in  ammonia,  forming 
Ag(NH3)2Cl;  reprecipitated  by  nitric  acid  as  AgCl.  Silver 
chloride  in  the  light  changes  from  white  to  lavender  and  finally 
to  black.1 


1  The  change  caused  by  sunlight  is  not  fully  understood.     Some  chlorine  is  given  off. 
Silver  chloride  precipitated  in  the  dark  has  slightly  different  properties  from  that  precipi- 


•:, 


METAL  ANALYSIS  6  1 

'precipitates  black  silver  sulphide,  Ag2S  ;.  s^uble  in 
hot  cftlute  nitric  acid. 

3.  NH4OH  precipitates  white  silver  hydroxide,  AgOH,  which 
rapidly  changes  to  brown  silver  oxide,  Ag2O  ;  soluble  in  excess 
of  reagent,  forming  Ag(NH3)2OH  and  r   whNH3)2]2O. 

4.  NaOH  or  KOH  precipitates  j  upon  corner  oxide,  Ag2O  ; 
soluble  in  ammonia  or  dilute  nitric  aof   plumbc 

5.  (NH4)2S  precipitates  black  Ag2S.  '  See  above. 

6.  Na2CO3  precipitates  white  silver  carbonate,   Ag2CO3,  or 
yellow  basic  carbonate. 

7.  K2CrO4  precipitates  from  neutral  solutions  bright  red  silver 
chromate,  Ag2CrO4  ;  soluble  in  nitric  acid. 

8.  KI    precipitates    pale   vellow   silver  iodide,   Agl  ;   easily 
soluble  in  excess  of  reagent^fc^oluble  in  ammonia  (!). 

9.  H2SO4    precipitates    from    concentrated    solutions    white 
silver  sulphate,  Ag2SO4. 

10.  KCN    precipitates   white    silver  cyanide,  AgCN  ;  easily 
soluble  in  excess,  forming 


1  1.    Na2HPO4  precipitates  yellow  silver  aphosphate,  Ag^PO^  ; 
soluble  in  dilute  nitric  acid,  and  ammonia. 

12.  Copper  and  some  other  metals  precipitate  metallic  silver. 

MERCURY  (ous)  :    Solution  for  Reactions,  Mercurous  Nitrate. 

13.  HC1  precipitates  white  mercurous  chloride,   HgCl;    in- 
soluble in  dilute  acids  ;  changed  by  ammonia  to  a  black  mixture 
of  Hg  and  HgNH2Cl.  % 

14.  H2S  precipitates  a  black  mixture  of  mercury,  Hg,  and 
mercuric  sulphide,  HgS.     K2S  in  presence  of  KOH  dissolves 
the  HgS,  leaving  Hg.     HgS  +  Hg  is  insoluble  in  nitric  acid, 
but  soluble  in  aqua  regia. 

tated  in  the  light.  Silver  chloride  exposed  to  the  light  is  acted  upon  quickly  by  mild  reduc- 
ing agent|j  alkaline  pyrogallate,  hydroquinone,  etc.,  depositing  metallic  silver,  while  that 
not  Exposed  to  light  is  not  so  rapidly  affected.  Dry  plate  photography  is  based  upon 
this  difference. 


to 


62  QUALITATIVE  ANALYSIS 

4 

15.  J§OH    or   NaOH    precipitates    a   black  fnixti! 
Hg2O,  and  HgO  ;  soluble  in  nitric  acid. 

1 6.  Na2CO3    precipitates     light     yellow    _^>asic    carbonates, 
changing  to  gray  bec^^^of  decomposition  into  HgO,  Hg,  and 
CO2.  *amn. 

17.  (NH4)2S  p  used  to  t5  a  black  mixture  of  Hg  and  HgS. 
See  14  above.       ?oing  exa- 

1 8.  K2CrO4  preCT^fEates  red  mercurous  chromate,  Hg2CrO4; 
insoluble  in  KOH. 

19.  NH4OH     precipitates     a    black    mixture    of    Hg    and 
HgNH2NO3;  soluble  in  aqua  regia. 

20.  KI  precipitates  yellowish  green  mercurous  iodide,  Hgl. 

21.  H2SO4  precipitates  white  mercurous  sulphate,  Hg2SO4; 


somewhat  soluble  in  water. 

r|^>i 


22.  KCN  produces^P|ray  pr|^>itate  of  Hg  and  Hg(CN)2.    * 

23.  SnCl2  reduces  mercurous  compounds  tu    metallic    mer- 
cury, Hg.  % 

24.  Copper  precipitates  metallic  mercury. 


Solution  for  Reactions,  Lead  Ac0ate  or  Nitrate. 

25.  HC1    precipitates,    incompletely,    lead    chloride,    PbCl2  ; 
soluble  in  boiling  water,  crystallizing   from  this  solution  upon 
cooling;  converted  into  white  basic  lead  chloride,  Pb(OH)Cl, 
by  ammonium  hydroxide. 

26.  H2S  precipitates  black  lead  sulphide  PbS,  even  in  the 
filtrate  frffi  the  lead  chloride  precipitation  ;  insoluble  in  yellow 
ammonium  sulphide  ;  soluble  in  warm  dilute  nitric  acid,  forming 
lead  nitrate,  Pb(NO3)2,  with  the  evolution  of  H2S.     Part  of  the 
H2S  is  oxidized  to  sulphur  jwith  the  reduction  of  the  nitric  acid 
to  the  oxides  of  nitrogen.     Hot  concentrated  nitric  acid  oxidizes 
lead  sulphide  to  white  lead  sulphate,  PbSO4. 

27.  NH4OH      precipitates      white      basic      lead     ciitrate, 
Pb3(OH)O2  •  NO3,  soluble  mJ4NO3.     With  lead  acetate  solu- 


in  H 


• 


**. 


63 

P  ordinary  strength,  excess  of  ammonium  hydroxide  (free 
from-  carbonate)  gives  no  precipitate. 

28.  (NH^S  precipitaTO  black  lead  sulphide,  PbS.  *See 
above. 

,29.    KOH    or    NaOH    precipitate  _  tehite    lead   hydroxide, 
Pb(OHA  or  basic  salts,  depending  upon  conditions  ;  soluble  in 
excess  of  reagent,  forming  salts   of  plumbous   acid,  Na2PbO/ 
and  K2PbO2  ]\  soluble  also  in  nitric  acid.J 

30.  (NH4)2CO3  precipitates  white  basic  lead  carbonates; 
soluble  in  strong  solution  x)f  potassium  or  sodium  hydroxides 
or  nitric  afcid.  » 


31.  NagCO^precipitetes  tne  same. 

32.  K2CrO4    precipitates    yellow    lead    chromate,    PbCrO4, 
"  chiome  yellow  "  ;    solubldjjn  sodium  hydroxide  ;    soluble  Vith 
difficulty  in  nitric  acid. 

33.  KI  precipitates  yellow  lead  iodide,  PbI2  ;  soluble  in 
watj^,  from  which  it  crystajil^s  upon  cooling  in  shining  gold 
yellow  scales. 

34.  H2SO4   or   a    soluble    sulphate    precipitates   white    lead 
sulphate,  PbSO4;    soluble  in  alkaline  ammonium  tartrate,  am- 
monium acetate,  or  sodium  hydroxide.     Lead  sulphate  is  less 
soluble  in  water  containing   alcohol   or   dilute   sulphuric   acid 
(why  ?)  than  in  water  alone.     Soluble  in  concentrated  H2SO4. 

35.  KCN   precipitates  white  lead  cyanide,  soluble   in  very 
large  excess  of  the  reagent.     Reprecipitated  upon  boiling. 

36.  SnCl2  produces  no  precipitate. 

37.  Zn  precipitates  metallic  lead  in  crystalline  form. 

Analysis,  Group  I 

Make  a  mixture  of  the  salts  of  the  metals  of  this  group,  and 
separate  by  the  following  scheme  of  analysis  : 

Dissolve  in  water  with  the  addition  of  a  few  drops  of  nitric 
acid  if  necessary. 


.- 


64  QUALITATIVE  ANALYSIS 

The  solution  is  tfeated  with  cold  dilute  hydrochloric  IK  as 
long  as  a  precipitate  is  formed.  Filter  and  wash  twice  with  cold 
water.  The  precipitate  contains  the^members  of  this  group. 

The  filtrate  may  contain  one  or  more  metals  of  Groups  II  to 
V,  and  should  be  saved  fiij^n  the  other  groups  are  to  be  analyzed. 

(A)  The  precipitates  washed  with  hot  water  several  times. 
The  residue  may  contain  mercury,  silver,  and  lead  (from  incom- 
plete washing).     The  filtrate  contains  lead  chloride. 

(B)  Divide  the  filtrate  into  several  portions  and  apply  the 
following  tests : 

Treat  with  H2SO4 ;  a  white  precipitate  indicates  lead. 
Treat  with  H2S  ;  a  black  precipitate  indicates  lead. 
Treat  with  K2CrO4 ;  a  yellow  precipitate   soluble  in   NaOH 
indicates  lead. 

Wash  with  hot  water  until  the  filtrate  no  longer  gives  a  test 
•  lead. 

(C)  To  the  precipitate  on  the  filter  add  ammonia ;   a  black 
residue  shows  the  presence  of  mercury. 


me 

€: 


actions  I 


E.    A  white  residue  insoluble  in  NH4OH  may  be  disregarded.     (Why  ?     See  re- 
forlead.) 

To  the  filtrate  add  an  excess  of  HNO3;  a  white  precipitate 
shows  the  presence  of  silver. 

After  having  completed  the  analysis  of  the  mixture,  proceed 
to  the  analysis  of  unknowns  Nos.  I  and  2.  These  being  reported, 
prepare  and  record  the  following  exercises : 

EXERCISES 

I.  Tabulate  the  reactions  which  are  common  to  all  the  mem- 
bers of  this  group. 

II.  Devise  another  method  for  the  analysis  of  the  members 
of  this  group. 

III.  Make  a  list  of  the  important  ores  of  lead,  silver,  and 
mercury.     Define  the  term  "ore." 


METAL  ANALYSIS  65 

What  metals  may  be  used  to  precipitate  the  ions  of  this 
group?     (See  Electromotive  Series.) 

V.  What  are  the  chief  uses  of  silver  and  lead  as  metals  ? 

VI.  What  is  white  lead,  and  how  is  it  made  ?     What  is  lunar 
caustic,  and  how  is  it  used  ? 

VII.  What  is  the  nature  of  the  compound  which  is  formed 
by  ammonium  hydroxide  acting  on  silver  salts  ?     Is  silver  am- 
photeric?      What   component   of   the    solution    of    ammonium 
hydroxide  effects  the  reaction? 

VIII.  Explain  the  formation  of   basic  salts  and  list  those 
which  are  formed  in  -the  preceding  reactions.     Define  the  term 
"basic  salt." 

Group  II 
Hydrogen  Sulphide  Group 

Mercury  (ic\  Lead,  Bismuth,  Copper,  Cadmium^  Arsenic, 
Antimony,  and  Tin 

GENERAL  STATEMENT.  The  sulphides  of  this  group 
are  insoluble  in  dilute  acids.  The  group  is  divided  into  two  sub- 
groups :  Subgroup  A,  those  metals  whose  sulphides  are  insoluble 
in  dilute  yellow  ammonium  sulphide ;  Subgroup  B,  those  metals 
whose  sulphides  are  soluble  in  dilute  yellow  ammonium  sulphide/ 

Subgroup  A 

Mercury  (ic),  Lead,  Bismuth,  Copper,  and  Cadmium 
MERCURY  (ic) :   Solution  for  Reactions,  Mercuric  Chloride. 

38.    HC1  produces  no  precipitate. 

39u  H2S,  slowly  added,  forms,  first  a  white  precipitate  (?); 
soluble  in  acids  and  in  excess  of  the  mercuric  salt.  By  further 
addition  of  H2S,  the  precipitate  becomes  orange-yellow,  then 
brown,  and  finally  black  mercuric  sulphide  is  produced ;  insolu- 


66  QUALITATIVE  ANALYSIS 

+ 
ble  in  ammonium  sulphide  or  hot  nitric  acid ;  soluble  in  aqua 

regia. 

NOTE.    The  white  precipitate  is  more  easily  secured  by  using  mercuric  nitrate. 

40.  NH4OH    precipitates    white    mercuric    amido    chloride, 
HgNH2Cl;  easily  soluble  in  HC1;  sparingly  soluble  in  strong 
ammonium  hydroxide. 

41.  (NH4)2S  precipitates  HgS.     The  same  changes  in  color 
as  in  (39)  may  be  noted  by  careful  addition  of  the  reagent. 

42.  KOH  or  NaOH   precipitates  from  cold  solutions,  first, 
reddish  brown  basic  salts,  which  change  to  orange-yellow  mer- 
curic oxide,  HgO,  when  the  reagent  is  added  in  excess.     When 
heated,  the  yellow  oxide  changes  to  red  HgO. 

In  the  presence  of  ammonium  salts,  the  white  "precipitate"  is 
formed.     See  40.     (Explain.) 

43.  (NH4)2CO3  acts  like  ammonium  hydroxide. 

44.  Na2CO3   precipitates,  first,  a  red-brown   basic   mercuric 
carbonate.     With  excess  of  reagents  and  heat,  it  is  converted 
into  yellow  mercuric  oxide,  HgO.     The  basic  salt  formed  with 
HgCl2  is  an  oxychloride,  HgCl2(HgO)2.  3.  or  4. 

45.  K2CrO4      precipitates      an      orange      basic     chromate, 
Hg2(OH)2Cr04. 

46.  KI  precipitates  yellow  mercuric  iodide,  HgI2  ;  soluble  in 
concentrated   HNO3  and  in   HC1 ;    soluble  in  solution  of  the 
iodides  of  the  more   positive   metals,  forming   double  iodides, 
as  K2HgI4. 

47.  KCN  precipitates  from  concentrated  solutions  white  mer- 
curic cyanide,  Hg(CN)2;  fairly  soluble  in  water;  easily  soluble 
in  KCN,  forming  a  double  salt,  Hg(CN)2  .  2  KCN.     Hg(CN)2  is 
not  acted  upon  by  the  hydroxides  of  the  alkali  metals.     (Ex- 
plain.    See  Appendix,  p.  177.) 

48.  SnCl2  precipitates,  in  the  cold,  white  mercurous  chloride, 
HgCl,  having  a  peculiar  silky  appearance  when  agitated.     If 


METAL  ANALYSIS  67 

excess  of  the  reagent  is  added,  the  mercurous  chloride  is  reduced 
more  or  less  completely  to  gray  or  black  metallic  mercury. 

49.  H2SO4  produces  no  precipitate. 

50.  Fused  with  Na2CO3  in  a  small  tube,  metallic  mercury  is 
deposited  on  the  walls  of  the  tube.     The  mercury  salt  should  be 
mixed  with  five  or  six  times  its  weight  of  Na^Og  and  heating 
begun  slowly  to  drive  out  any  moisture  which  may  be  present. 
This  may  be  wiped  out  with  a  piece  of  filter  paper. 

LEAD  :    Solution  for  Reactions,  Lead  Nitrate  or  Acetate. 

51.  The  reactions  for  lead  are  given  on  p.  62. 
BISMUTH:    Solution  for  Reactions,  Bismuth  Chloride. 

52.  Water  precipitates   white  bismuth  oxychloride,  BiOCl ; 
soluble   in    hydrochloric    acid,    but   insoluble    in    tartaric   acid. 
(See  Sb.)     Converted  by  H2S  to  Bi2S3. 

53.  HC1  produces  no  precipitate. 

54.  H2S  precipitates  black  bismuth  sulphide,  Bi2S3  ;  insoluble 
in  dilute  acids  and  alkalies  and  in  ammonium  sulphide,  soluble 

in  modej»n?ly  concentrated  nitric  acid. 
^•fc*   *i 
J<JH  precipitates  white  bismuthyl  hydroxide,  BiOOH  ; 

ible  in  excess  ;  soluble  in  dilute  acids ;  converted  by  boil- 
Finto  yellowish  white  bismuth  oxide,  Bi2O3;  action  facilitated 
bj^addition  of  strong  alkali. 

|6.    (NH^S  precipitates   bfack   bismuth  sulphide.     See  54 

NaOH  or  KOH  precipitates  same  as  ammonium  hydrox- 

ide>  jr 

58.  (NH4)2CQg   Sf^Na3CO3    precipitates    white    bismuthyl 
carbonate,  (BiO)2CO3 ;  soluble  in  dilute  acids. 

59.  K2CrO4    precipitates   yellow    basic   bismuth    chromate, 
Bi2O(CrO4)2  ;  insoluble  in  sodium  hydroxide  ;  soluble  in  HNO3. 

60.  KI    produces,  in   solutions  not  too  strongly   acid,  dark 
brown  bismuth  iodide,  BiI3.     If  the  precipitate  does  not  form 


68  QUALITATIVE  ANALYSIS 

after  adding  the  reagent,  make  slightly  less  acid  with  NaOH 
until  it  appears  ;  soluble  in  excess  of  KI  and  HC1 ;  not  soluble  in 
dilute  nitric  acid. 

61.  H2SO4  produces  no  precipitate. 

62.  SnCl2  produces  no  precipitate. 

63.  Zn  or  Fe  precipitates  spongy  bismuth. 

COPPER  :    Solution  for  Reactions,  Copper  Sulphate. 

64.  HC1  produces  no  precipitate. 

65.  H2S  precipitates  black  copper  sulphide,  CuS  ;  soluble  in 
nitric  acid  and  potassium  cyanide ;  insoluble  in  yellow  ammonium 
sulphide. 

66.  NH4OH     precipitates     pale    blue     copper    hydroxide, 
Cu(OH)2;  soluble  in  excess  of  reagent,  forming  a  deep  blue 
solution  ;  decolorized  by  KCN. 

67.  (NH4)2S   precipitates   black   copper   sulphide.     See   65 
above. 

68.  KOH  or  NaOH  precipitates  pale  blue  copper  hydroxide, 
Cu(OH)2 ;  changed  by  boiling  with  excess  of  alkali  to  black 
copper  oxide,  CuO.  *** 

69.  Na2CO3  precipitates  greenish  blue  basic  copper  carbonate, 
Cu2(OH)2CO3;    converted  by  foiling  to   tlack    copper   oxide, 
CuO. 

70.  K2CrO4  precipitates  reddish  brown  basic  copper  chromat)£ ; 
somewhat  soluble  in  water ;  soluble  in  acids. 

71.  KI  precipitates  from  concentrated  solutions  white  cuprous 
iodide,  Cu2I2,  with  the  liberation  of  iodine. 

72.  KCN    precipitates     yellowish    green     copper     cyanide, 
Cu(CN)2  ;  soluble  in  excess  of  reagent,  producing  a  double  salt, 
K2Cu(CN)4.     Not  precipitated  by  H2S.     (See  p.  39.) 

73.  H2SO4  produces  no  precipitate. 

74.  K4Fe(CN)6  precipitates  reddish  brown  copper  ferrocyan- 
ide,  Cu2Fe  (CN)6;  insoluble  in  acids  ;  decomposed  by  alkalies. 


METAL  ANALYSIS  69 

75.  SnC^  produces  no  precipitate. 

76.  Zn  and  Fe  precipitate  metallic  copper. 
CADMIUM  :    Solution  for  Reactions,  Cadmium  Sulphate. 

77.  HC1  produces  no  precipitate. 

78.  H2S    precipitates   cadmium    sulphide,   CdS,   varying   in 
color  from  bright  yellow  to  orange,  according  to  conditions ; 
easily  soluble  in  HC1,  hot  H2SO4,  or  hot  HNO3 ;    insoluble  in 
yellow  ammonium  sulphide  and  potassium  cyanide. 

79.  NH4OH  precipitates  white  cadmium  hydroxide,  Cd(OH)2 ; 
soluble  in  excess.     If  the  ammoniacal  solution  is  treated  with 
potassium  cyanide,  a  soluble  double  salt,  K2Cd(CN)4,  is  formed 
from  which  the  cadmium  may  be  precipitated  by    H2S.     See 
reaction  72. 

80.  (NH^S  precipitates   yellow   cadmium    sulphide,   CdS. 
See  78. 

8 1.  KOH  or  NaOH  precipitates  white  cadmium  hydroxide, 
Cd(OH)2;  insoluble  in  excess;  soluble  in  acids. 

82.  K^CrC^  precipitates,  from  concentrated  solutions   only, 
yellow    cadmium    chromate,  CdCrO4 ;    soluble   on    addition    of 
water. 

83.  KI  produces  no  precipitate. 

84.  H2SO4  produces  no  precipitate. 

85.  KCN    precipitates   white   cadmium   cyanide,  Cd(CN)2; 
soluble  in  excess,  forming  K^CdCCN^;  reprecipitated  by  H2S 
as  cadmium  sulphide. 

86.  SnCl2  produces  no  precipitate. 

Analysis,  Group  II,  Subgroup  A 

Make  a  mixture  of  the  salts  of  the  metals  of  this  subgroup ; 
dissolve  in  water  with  the  aid  of  hydrochloric  acid,  and  separate 
and  identify  by  the  following  scheme  of  analysis : 

The  solution  to  which  hydrochloric  acid  is  added  is  boiled 
nearly  to  dryness  to  decompose  any  nitric  acid  which  may  be 


70  QUALITATIVE  ANALYSIS 

present.  It  is  then  diluted  with  water,  and  a  small  amount  of 
hydrochloric  acid  is  added  if  necessary l  to  dissolve  any  residue 
which  may  be  present. 

NOTE.  As  nitric  acid  decomposes  hydrogen  sulphide,  the  solution  should  be  boiled 
nearly  to  dryness  with  frequent  additions  of  small  quantities  of  hydrochloric  acid  to 
remove  the  nitric  acid,  unless  it  is  known  that  no  nitric  acid  or  nitrates  are  present. 

The  solution  is  now  heated  to  boiling  and  hydrogen  sulphide 
run  inyslowly2  until  saturated.  The  solution  is  then  cooled, 
diluted  with  water,  and  more  hydrogen  sulphide  run  in. 

The  precipitate,  which  contains  the  metals  of  Group  II,  is 
filtered  off  and  washed  with  hot  water  until  a  few  drops  of  the 
filtrate  impart  only 'slight  turbidity  to  silver  nitrate  solution. 

The  filtrate  contains  the  metals  of  Group  III  to  V,  and  should 
be  saved  if  these  are  to  be  sought. 

(A)  The     precipitate    is    warmed    with    dilute     ammonium 
sulphide   and  filtered.     The  filtrate  may  contain    members   of 
Subgroup  B,  and  should  be  saved  if  the  members  of  this  group 
are  sought. 

The  precipitate  insoluble  in  ammonium  sulphide  contains 
all  the  members  of  Subgroup  A.  Analyzed  by  (B)  and 
following : 3 

(B)  The  precipitate  is  boiled  with  dilute  nitric  acid  for  a  few 
minutes,  filtered,    and   washed.     The  precipitate  may  contain 
mercuric   sulphide  (black),   lead   sulphate   (white),  or   sulphur 
(yellow,  or  black,  and  floating). 

Test  for  mercury  by  adding  aqua  regia,  boiling,  evaporating 
to  expel  nitric  acid,  diluting  with  water  and  hydrochloric  acid, 
and  adding  stannous  chloride.  See  reaction  48. 

1  The  precipitation   of  this  group  is   accomplished  with  best  results  if  the  solution 
contains  about  one  part  of  HC1  in  ten  of  H2O. 

2  It  is  suggested  that  the  solution  be  placed  in  an  Erlenmeyer  flask,  fitted  with  a  one- 
holed  stopper,  through  which  runs  a  glass  tube  nearly  to  the  bottom  of  the  flask.     Attach 
to  the  H2S  generator,  loosen  the  stopper,  and  pass  in  H2S  until  the  air  is  replaced,  then 
close  tightly  and  run  in  H2S,  with  frequent  shaking  of  the  flask,  until  action  ceases. 

3  If  Subgroup  B  is  known  to  be  absent,  as  in  the  preliminary  separation,  (A)   may 
be  omitted. 


METAL   ANALYSIS  71 

(C)  The  filtrate  from  (B)  is  treated  with  2  c.c.  of  sulphuric 
acid  and  evaporated  until  white  fumes  of  sulphur  trioxide  are 
given  off.     Cool,  add  water,  and  filter  if  necessary  (Filtrate  D). 
A  white  precipitate  indicates  lead.    Add  concentrated  ammonium 
acetate1  and  a  little  acetic  acid  to  the  precipitated  lead  sulphate, 
warm,  and  filter.     To  the  filtrate  add  potassium  chromate.     A 
yellow  precipitate  soluble  in  sodium  hydroxide  shows  the  pres- 
ence of  lead. 

(D)  The  filtrate  from  the  sulphuric  acid  precipitate  is  treated 
with  an  excess  of  ammonia.     Filter  if  necessary.     To  the  precip- 
itate add  a  few  drops  of  cone.  HC1;  evaporate  to  small  volume 
and  pour  into  about  500  c.c.  of  water.     A  white  precipitate  shows 
the  presence  of  bismuth. 

(E)  The  filtrate  from  (D)  may  contain  copper  and  cadmium. 
If  the  filtrate  is  blue,  copper  is  present.     If  faintly  blue,  confirm 
the  presence  of  copper,  in  a  small  portion,  by  74. 

(F)  If  copper  is  present,  decolorize  with  potassium  cyanide 
and  treat  with  hydrogen  sulphide.     A  yellow  precipitate  shows 
the  presence  of  cadmium.2 

After  having  completed  the  analysis  of  the  mixture,  proceed 
to  the  analysis  of  unknowns  Nos.  3  and  4.  These  being  reported, 
prepare  and  record  the  following : 

EXERCISES 

I.  What  reactions  are  common  to  all  the  metals  of  this  sub- 
group ? 

II.  Explain  the  presence  of  sulphur  and  lead  sulphate  in  (B). 

III.  If  the  filtrate  from  the  mercuric  sulphide  in  (B)  were 
treated  at  once  with  ammonium  hydroxide,  what  might  the  pre- 
cipitate contain  ? 

1  2  PbS04  +  2  NH4C2H302  ->  Pb(XH4)2(S04)2  +  Pb(C2H3O2)2. 

2  If  the  precipitate  is  not  yellow,  it  may  be  redissolved  in  dilute  hydrochloric  acid, 
boiled,  made  alkaline  with  an  excess  of  ammonia,  filtered,  and  reprecipitated  with  hydrogen 
sulphide. 


72  QUALITATIVE  ANALYSIS 

IV.  What   indications  of   the    presence  or  absence  of  cer- 
tain metals  are  given  by  the  color  of  the  hydrogen  sulphide 
precipitate  ? 

V.  What  are  the  ions  of  hydrogen  sulphide,  which  are  most 
numerous  ?     Why  are  sulphides  precipitated  instead  of  hydro- 
sulphides  ?     Cf.  Smith,  General  Chemistry,  p.  374. 

VI.  What   are   the  ions    present   in   a   solution  of   copper 
hydroxide  ?     Does  copper  act  as  an  acid-forming  element  ? 

VII.  Make  a  table  showing  mineralogical  names  and  formulas 
of  the  important  ores  of  each  metal  of  the  group. 

VIII.  Give  the  chemical  reactions  for  at  least  one  method  of 
preparation  of  each  of  these  metals. 

IX.  What  are  the  uses  of  calomel  and  corrosive  sublimate, 
and  how  are  they  distinguished  from  each  other  ?    Smith,  General 
Chemistry,  p.  654. 

X.  What  is  "  red  precipitate  "  ?     "  White  precipitate  "  ? 

XL     What   is    Wood's    metal  ?     Rose's    metal  ?     What    are 
amalgams  ? 

XII.  What  are  the  uses  of  bismuth  subnitrate  ?   of  copper 
sulphate  ?  of  cadmium  sulphide  ? 

XIII.  What  are  the  chief  alloys  of  copper  ? 

XIV.  Why  is  copper  not  precipitated  by  hydrogen  sulphide 
from  the  solution  containing  an  excess  of  KCN,  while  cadmium 
is  precipitated  ?     Why  is  the  blue  solution  decolorized  by  KCN  ? 

Subgroup  B 

Metals  whose  Sulphides  are  Soluble  in  Dilute  Yellow  Ammonium  Sulphide 
Arsenic,  Antimony,  and  Tin 

ARSENIC  (ous) :    Solution  for  Reactions,  Sodium  Arsenite  or 
Arse  nous  Chloride.^- 

87.  HC1  produces  no  precipitate. 

88.  H2S  precipitates,  from  hydrochloric  acid  solution,  yellow 

1  Arsenous  chloride  is  prepared  by  dissolving  arsenous  oxide  in  hydrochloric  acid. 


METAL  ANALYSIS  73 

arsenous  sulphide,  As2S3;  soluble  in  yellow  ammonium  sul- 
phide, forming  sulpho  salts,  such  as  (NH4)3AsS4  ;  reprecipitated 
as  As2S5  and  As2S3  by  dilute  HC1 ;  soluble  in  sodium  hydroxide, 
ammonium  carbonate,  hot  nitric  acid,  and  nascent  chlorine, 
nearly  insoluble  in  warm  concentrated  hydrochloric  acid. 

89.  NH4OH,  KOH,  or  NaOH  produces  no  precipitate. 

90.  (NH^S  precipitates  As2S3  from  acid  solutions  ;  soluble 
in  excess.     See  88  above. 

91.  (NH4)2CO3  or  Na^COg  produces  no  precipitate. 

92.  AgNO3  precipitates  from  neutral  solutions  silver  arsenite, 
Ag3AsO3 ;  readily  soluble  in  dilute  acids,  ammonium  hydroxide, 
or  ammonium  salts. 

93.  CuSO4  in  neutral  solution  precipitates  "  Scheele's  green," 
CuHAsO3;  soluble  in  ammonium  hydroxide  and  dilute  acids. 

94.  A  bright  strip  of  copper  placed  in  a  solution  of  arsenic 
(ous),  made  strongly  acid  with  HC1  and  boiled,  gives  a  gray 
film  of  arsenic  on  the  copper.     "  Reinsch's  test." 

95.  Oxidizing  agents,  as  potassium  chromate  and  potassium 
permanganate,   convert   arsenous    compounds  to  arsenic   com- 
pounds. 

96.  Nascent    hydrogen1    reduces    arsenic    compounds   to    a 
colorless  poisonous  gas,  arsine,  AsH3,  decomposed  by  heating 

1  Marsh's  Test  for  Arsenic.  —  A  flask  containing  arsenic  free  zinc  is  fitted  with  a  funnel 
tube  and  a  delivery  tube  of  glass  drawn  to  a  small  point  at  the  end.  Hydrochloric  acid  is 
added  to  the  zinc  in  the  flask  through  the  funnel  tube  until  hydrogen  is  freely  evolved. 
As  soon  as  all  of  the  air  is  out  of  the  apparatus  (test)  (see  Instructor),  light  and  test  for 
arsenic  or  allied  metals  by  holding  a  cold  porcelain  dish  in  the  flame.  If  no  spot  is 
formed  on  the  dish,  the  apparatus  is  ready  for  carrying  out  the  test  for  arsenic.  Add 
through  the  funnel  tube  a  hydrochloric  acid  solution  of  an  arsenic  compound.  Heat  the 
delivery  tube  near  the  tip.  Note  the  black  metallic  mirror  of  arsenic  formed.  Write 
equation.  Light  the  issuing  gas  at  the  tip  of  the  delivery  tube  and  hold  a  cold  porcelain 
dish  in  the  flame.  Note  the  spot  ot  black  metallic  arsenic.  Get  several  spots  on  the 
dish  and  save  for  future  test  No.  98.  Put  out  the  flame  and  run  the  gas  into  a  solution  of 
silver  nitrate  in  a  test  tube.  A  precipitate  of  black  metallic  silver  is  formed.  Add  HC1  to 
the  solution  in  the  test  tube  until  a  precipitate  no  longer  forms.  Filter  and  test  the  filtrate 
for  arsenic  by  H2S. 


74  QUALITATIVE  ANALYSIS 

to  metallic  arsenic  and  hydrogen ;  reduces  a  solution  of  silver 
nitrate  to  metallic  silver,  changing  at  the  same  time  to  arsenous 
acid,  H3AsO3,  which  may  be  precipitated  as  arsenic  trisulphide 
with  hydrogen  sulphide. 

97.  KCN    and   Na^Og  fused  with    arsenic   trioxide  or  tri- 
sulphide in  a  tube  produce  deposits  of  metallic  arsenic  on  the 
cold  parts  of   the  tube  as  a  shining  black  mirror,  the  KCN 
changing  to  KCNO  or  KCNS.     The  moisture  should  first  be 
removed  in  this  experiment  as  directed  in  50. 

98.  Arsenic  spots1  dissolve  in  sodium  or  calcium  hypochlo- 
rite  and  in  hot  nitric  acid. 

ARSENIC:  Solution  for  Reactions,  Sodium  Ar senate. 

99.  HC1  produces  no  precipitate. 

100.  H2S  precipitates   yellow  arsenic  sulphide,  or  arsenous 
sulphide,  As2S3,  and  sulphur  depending  upon  conditions.     The 
second  reaction  takes  place  more  freely  when  hydrogen  sulphide 
is  passed  slowly  through  a  hot  solution.     Both  are  soluble  in 
ammonium  sulphide.     See  88. 

101.  NH4OH  produces  no  precipitate. 

102.  KOH  or  NaOH  produces  no  precipitate. 

103.  (NH^S    precipitates   yellow    arsenic   sulphide,    As2S3, 
from  solutions  acidified  with  hydrochloric  acid. 

104.  (NH4)2CO3  or  Na2Co3  produces  no  precipitate. 

105.  AgNO3  precipitates  from  neutral  solutions  reddish  brown 
silver  arsenate,   Ag3AsO4 ;    readily    soluble  in    dilute  acids  or 
ammonia. 

1 06.  CuSO4  precipitates    greenish    blue   acid    arsenates    as 
CuHAsO4;   soluble  in  dilute  acids  and  ammonia. 

107.  Magnesium    mixture2    precipitates     white    magnesium 
ammonium  arsenate,  MgNH4AsO4;  easily  soluble  in  acids. 

1  The  spots  produced  by  burning  AsH3  when  a  cold  dish  is  held  in  the  flame  are  called 
arsenic  spots. 

2  For  preparation,  see  Appendix. 


METAL   ANALYSIS  75 

108.  Nascent  hydrogen  reduces  arsenic  compounds  to  arsine, 
AsH3.     See  96. 

109.  Reducing   agents   FeSO4,    Na^Og,   etc.,    heated    with 
H2SO4,  reduce  arsenic  acid  to  arsenous  acid,  H3AsO3. 

no.    Heated  with  charcoal  in  a  closed  tube,  a  mirror  of  metallic 
arsenic  is  formed  on  the  walls  of  the  tube, 
in.    Arsenic  spots.     See  98. 

ANTIMONY  :  Solution  for  Reactions,  Antimony  Chloride. 

112.  HC1  produces  no  precipitate. 

113.  H2S  precipitates  orange-red  antimony  sulphide,  Sb2S3; 
soluble  in  ammonium  sulphide,  forming  (NH4)3SbS4  (see  As); 
soluble  in  sodium  hydroxide  and  in  warm  concentrated  hydro- 
chloric acid. 

114.  NH4OH      precipitates     white      antimony      hydroxide, 
Sb(OH)3;  soluble  in  dilute  hydrochloric  acid  or  sodium  hydrox- 
ide. 

115.  NaOH  or  KOH  precipitates  the  same ;  soluble  in  excess 
of  reagent  or  in  hydrochloric  acid. 

1 1 6.  (NH^S    precipitates   from   acid    solutions   orange-red 
antimony   trisulphide,   Sb2S3;    soluble  in   ammonium  sulphide. 
See  113. 

117.  (NH^COg    or    Na2CO3    precipitates    white   antimony 
hydroxide,  Sb(OH)3;  soluble  upon  warming  in  large  excess  of 
sodium  carbonate,  but  not  in  ammonium  carbonate;   soluble  in 
hydrochloric  acid. 

1 1 8.  Water  precipitates  white  antimony  oxychloride,  SbOCl; 
soluble  in  acids ;  converted  by  H2S  into  Sb2S3. 

119.  Nascent  hydrogen  (see  Marsh's  test)  reduces  antimony 
compounds  to  stibine,  SbH3. 

1 20.  Antimony  spots  are  not  dissolved  in   hypochlorites,  but 
are  by  concentrated  HNO3. 


76  QUALITATIVE  ANALYSIS 

1 2 1.1  Tin2  in  presence  of  hydrochloric  acid  and  platinum  foil 
precipitates  metallic  antimony  as  a  dark  brown  powder  adher- 
ing to  the  platinum  ;  concentrated  nitric  acid  causes  the  stain  to 
disappear;  forming  meta-antimonic  acid,  HSbO3.  If  the  liquid 
is  touched  with  a  glass  rod  dipped  in  an  ammoniacal  solution  of 
silver  nitrate,  white  silver  meta-antimonate,  AgSbO3,  is  formed. 

TIN  (STANNOUS)  :  Solution  for  Reactions,  Stannous  Chloride 

122.  HC1  produces  no  precipitate. 

123.  H2S  precipitates  dark  brown  stannous  sulphide,  SnS  ; 
soluble   in    yellow,    but   not  in  colorless,  ammonium  sulphide, - 
forming  ammonium  sulphostannate,  (NH4)2SnS3;  reprecipitated 
by  dilute  acids,  forming  yellow  stannic  sulphide,  SnS2. 

124.  NH4OH      precipitates      white      stannous      hydroxide, 
Sn(OH)2;  insoluble  in  excess;  soluble  in  dilute  HC1. 

125.  NaOH    or    KOH    precipitates    the    same;    soluble    in 
excess,  forming  stannites,  Sn(ONa)2  or  Sn(OK)2. 

126.  Na2CO3  precipitates  the  same  as  1254 

127.  (NH^S    precipitates    from  acid  solutions  dark  brown 
stannous  sulphide,  SnS.     See  123. 

128.  HgCl2   is   reduced  by  SnCL,  forming  white  HgCl,  or 
black    Hg   if    SnCl2   is   in  excess  ;  at  the  same  time  SnCl2  is 
oxidized  to  stannic  chloride,  SnCl4. 

129.  Zinc  in  hydrochloric  acid  solution  precipitates  metallic 
tin   as  a   spongy    mass    upon   the  zinc.     For  distinction  from 
antimony,  see  121. 

130.  Oxidizing  agents,  such  as  HNO3,  KC1O3,  etc.,  in  HC1, 
convert  stannous  salts  to  stannic  salts. 

1  To  perform  this  test  a  few  drops  of  the  solution  are  placed  upon  a  clean  platinum 
foil ;  dip  one  end  of  a  U-shaped  strip  of  tin  foil  in  the  liquid,  leaving  the  other  end  in  con- 
tact with  the  platinum  beyond  the  drop.    A  bright  silver  coin  may  be  used  instead  of 
platinum. 

2  Tin  is  used  rather  than  zinc,  because  the  latter  also  precipitates  tin,  which  may  be 
mistaken  for  antimony,  although  the  deposit  is  much  brighter. 


METAL  ANALYSIS  77 

TIN  (STANNIC)  :  Solution  for  Reactions,  Stannic  Chloride! 

131.  HC1  produces  no  precipitate. 

132.  H2S  precipitates   yellow    stannic   sulphide;    soluble   in 
warm  concentrated  HC1 ;  soluble  in  yellow  ammonium  sulphide, 
forming  (NH4)2SnS3.     See  123. 

133.  NH4OH  precipitates  white  metastannic  acid,  SnCXOH^. 

134.  NaOH  precipitates  the  same;  soluble  in  excess. 

135.  Na^Og  precipitates  the  same. 

136.  (NH^S    precipitates    yellow    stannic  sulphide,    SnS2; 
soluble  in  yellow  ammonium  sulphide. 

137.  HgCl2  produces  no  precipitate. 

138.  Zinc  precipitates  metallic  tin.     See  129. 

Analysis,  Group  II,  Subgroup  B 

Make  a  mixture  of  the  salts  of  the  metals  of  this  subgroup, 
and  separate  by  the  following  scheme  of  analysis ;  2 

(A)  The  filtrate3  from   (A),    Subgroup   A,  is   treated   with 
dilute  HC1  to  acid  reaction.     Filter  and  wash  with  hot  water. 
The  filtrate  is  discarded.     The   precipitate,  containing  all  the 
members  of  this  subgroup  present,  is  treated  with  concentrated 
hydrochloric  acid  and  warmed  for  a  few  minutes,  filtered,  and 
washed. 

(B)  The  residue  may  contain  arsenic  and  sulphur;  boil  with 
concentrated  HC1  and  a  few  crystals  of  KC1O3,  gradually  added. 
Take  out  the  floating  sulphur  with  a  glass  rod ;  make  alkaline 
with  ammonia,  and  test  the  solution  for  arsenic  with  magnesium 
mixture.     (See  other  tests.) 

1  The  stannic  chloride  used  for  these  preliminary  reactions  should  be  prepared  from 
stannous  chloride  by  treatment  with  HC1  and  KC1O3  and  heating  until  chlorine  is  no 
longer  evolved.      Metastannic  acid  gives  somewhat  different  reactions,  but  since  it  is  pre- 
cipitated by  H2S  as  stannic  sulphide,  it  is  considered  sufficient  to  use  the  stannic  salt  given 
above  for  these  preliminaries. 

2  Begin  the  analysis  for  Subgroup  B  as  given  for  analysis  of  Subgroup  A. 

3  If  Subgroup  A  is  known  to  be  absent,  the  solution  and  reprecipitation  with 
and  HC1  may  be  omitted  and  the  analysis  begun  with  the  H2S  precipitate. 


78  QUALITATIVE  ANALYSIS 

(C)  The  filtrate   from  (A)  may  contain   tin  and   antimony.1 
Add  an  excess  of  zinc,2  and  a  piece  of  platinum  foil,  taking 
care  that  the  platinum  touches  the  zinc. 

(D)  The  black  coating  on  the  platinum  is  thoroughly  washed 
with  hot  water  and  touched  with  a  drop  of  nitric  acid,  and  then 
a  drop  of  ammoniacal  silver  nitrate.     A  white  deposit  confirms 
the  presence  of  antimony. 

(E)  The  excess  of  zinc  with  the  coating  of  tin  is  dissolved  in 
HC1,  boiled  nearly  to  dryness  to  expel  excess  of  acid,  diluted, 
and  treated  with  mercuric  chloride.     A  white  or  gray  precipitate 
confirms  the  presence  of  tin.     If  present  in  small  quantities,  it 
is  best  to  boil ;  a  black  precipitate  of  Hg  after  standing  a  few 
minutes  confirms  the  presence  of  tin. 

Alternative  analysis.  The  filtrate  from  (A)  may  be  tested 
for  antimony  and  tin  without  the  use  of  platinum  as  follows: 
Boil  to  expel  hydrogen  sulphide  and  place  a  few  drops  on  a 
bright  silver  surface  (coin).  A  U-shaped  strip  of  tin  foil  is 
placed  so  that  one  end  touches  the  silver  beyond  the  drop. 
In  the  presence  of  antimony  a  dark  stain  appears  on  the 
silver. 

Another  portion  is  tested  for  tin  by  nearly  neutralizing  the 
acid,  and  adding  a  bit  of  granulated  zinc  so  that  some  zinc  re- 
mains after  action  has  ceased.  The  residue  of  zinc  with  the 
adhering  tin  is  dissolved  in  hydrochloric  acid  and  tested  as  in 
(E)  above. 

After  having  completed  the  analysis  of  the  mixture,  proceed 
to  the  analysis  of  unknown  Nos.  5  and  6.  These  being  reported, 
prepare  and  record  the  following : 

1  A  very  satisfactory  identification  of  antimony  is  carried  out  as  follows  :     Take  a  drop 
of  the  solution  containing  antimony  and  tin  and  place  on  a  platinum  foil;  place  a  U- 
shaped  strip  of  tin  with  one  end  in  and  the  other  outside  the  drop  in  contact  with  the 
platinum  foil.    A  black  precipitate  on  the  foil  is  conclusive  evidence  of  the  presence  of 
antimony. 

2  So  that  when  action  ceases,  there  will  be  some  zinc  left. 


METAL   ANALYSIS  79 

EXERCISES 

I.  Devise  another  method  for  the  separation  of  the  members 
of  this  subgroup,  if  in  solution  by  themselves. 

II.  Of  what  use  is  KC1O3  in  (B)?     What  is  ordinary  com- 
mercial "  arsenic  "  ? 

III.  Explain   why    SnS  is   not   dissolved  by   colorless  am- 
monium sulphide,  and  why  As2S3  and  Sb2S3  are  dissolved  only 
with  difficulty. 

IV.  What   are   the   chief   ores   of   arsenic,    antimony,   and 
tin? 

V.  Make   a   table  of   the  antimony,  arsenic,  and   arsenous 
acids,  showing  names  and  formulae. 

VI.  How  is  most  of  the  "  arsenic  "  of  commerce  obtained  ? 

VII.  What  is  "tin  plate"?     How   produced?     Banca   tin? 
Block  tin  ? 

VIII.  What  is  tartar  emetic  ?     How  made  ? 

Group  III 

The  Ammonium  Sulphide  Group 

Iron,   Chromium,  Aluminium,  Manganese,  Zinc,  Nickel,  and 

Cobalt 

GEN ERAL  ST ATEM ENT.  Iron,  chromium,  and  aluminium 
are  precipitated  as  hydroxides  by  ammonium  hydroxide  in  the 
presence  of  ammonium  chloride.  Manganese,  zinc,  nickel,  and 
cobalt  are  not  precipitated  unless  phosphoric  or  some  similar 
acid  is  present.  Ammonium  sulphide  converts  all  the  metals 
into  sulphides  except  chromium  and  aluminium,  which  come 
down  as  hydroxides.  In  the  presence  of  phosphates,  borates, 
oxalates,  silicates,  fluorides,  or  tartrates,  barium,  strontium,  cal- 
cium, and  magnesium  may  be  precipitated  in  this  group.  For 
test  for  the  presence  of  these  acids,  see 'acid  analysis. 


80  QUALITATIVE  ANALYSIS 

IRON  (FERROUS):     Solution  for  Reactions,  Ferrous  Sulphate. 

139.  HC1  produces  no  precipitate. 

140.  H2S  in  acid  solution  produces  no  precipitate./ 

141.  NH4OH  precipitates,  incompletely,  ferrous  hydroxide/' 
changing  slowly  to  reddish  brown  ferric  hydroxide,  Fe(OH)3 ; 
soluble  in  acids,  even  acetic  acid ;  treated  with  (NH4)2S  forms 
ferrous  sulphide,  FeS,  black. 

142.  (NH4)2S    precipitates    black    ferrous    sulphide,    FeS; 
soluble  in  HC1. 

143.  NaOH  or  KOH  precipitates  white  ferrous  hydroxide/ 
oxidizing  immediately  to  a  dirty  green,  to  black,  and  finally -to 
reddish  brown  ferric  hydroxide.     See  141.  - 

144.  (NH4)2CO3  precipitates  white  ferrous  carbonate,  FeCO3, 
which  quickly  oxidizes  and  darkens  in  color,  eventually  changing 
to  Fe(OH)3;  soluble  in  acids,  even  acetic  acid. 

145.  NagHPC^  precipitates  a  mixture  of  acid  ferrous  phos- 
phate, FeHPO4,    and  ferrous  phosphate,  Fe3(PO4)2,    white   to 

bluish   white.     By  the   addition   of   an  alkali   acetate,  ferrous 

'    rift 
phosphate  alone,  Fe3(PO4)2,  is  formed;  soluble  in  HC1  or  HNOg.1 

146.  K4Fe(CN)6    precipitates    white    potassium    ferroferro- 
cyanide,    K2FeFe(CN)6,  quickly   turning   blue  by  oxidation,  a 
small   amount    of    ferric^^i^cyanide,    Fe4(Fe(CN)6)3,    being 
formed. 

147.  K3Fe(CN)6  precipitates  dark  blue  ferrous  ferricyanide, 
Fe3(Fe(CN)6)2,  "  Turnbull's  blue";  insoluble  in  dilute  acids. 

148.  KCNS    produces    no   coloration    when    ferric    iron    is 
entirely  absent. 

149.  Boiled  with  HNO3,  a  ferric  salt  is  formed. 

150.  Chlorine  water  oxidizes  a  solution  of  ferrous  salt  to  a 
ferric  salt. 

151.  Barium    carbonate  suspended   in  water,    when    shaken 

1  For  this  and  the  two  following  tests  a  ferrous  salt  free  from  ferric  iron  must  be  used. 
See  appendix  for  preparation. 


METAL  ANALYSIS  8 1 

with  a  cold  neutral  or  slightly  acid  solution  of  a  ferrous  salt, 
produces  no  precipitate  containing  iron.  Why  ?  (See  Table 
of  Solubilities.)  Cf.  163  and  171. 

IRON  (FERRIC):  Solution  for  Reactions,  Ferric  Chloride. 

152.  HC1  produces  no  precipitate. 

153.  H2S  reduces  solutions  of  ferric  salt  to  the  ferrous  con- - 
dition,  sulphur  being  set  free.' 

154.  NH4OH    precipitates  reddish   brown  gelatinous    ferric 
hydroxide,  Fe(OH)3  ;  soluble  in  acids. 

155.  KOH  or  NaOH  precipitates  the  same. 

156.  (NH^S  precipitates  ferrous  sulphide,  FeS,  mixed  with 
free  sulphur. 

157.  (NHJ2CO3  or  Na2CO3  precipitates  Fe(OH)3. 

1 58.  Na2HPO4  precipitates  ferric  phosphate,  FePO4  ;  slightly 
soluble  in  acetic  acid;  readily  soluble  in  HC1,  HNO3,  and  H2SO4. 

159.  K4Fe(CN)6     precipitates     blue     ferric     ferrocyanide, 
Fe4(Fe(CN)6)3,  "  Prussian  blue ";  insoluble  in  dilute  inorganic 
acids. 

1 60.  K3Fe(CN)6    produces    no    precipitate,    but    imparts   a 
reddish  brown  color. 

161.  KCNS  produces  an  intense  red  color  of  ferric  sulpho- 
cyanate,  Fe(CNS)3.  +* 

162.  Most  reducing  agentrfas  H2S,  SO2,  and  SnCl2,  easily 
reduce  ferric  to  ferrous  salts. 

163.  BaCO3  suspended  in  water,  when  shaken  with  a  cold 
neutral  or  slightly  acid  solution  of  a  ferric  salt,  precipitates 
brown  ferric  hydroxide,  Fe(OH)3;    soluble  in   HC1,   and  may 
be  confirmed  by  161. 

CHROMIUM:  Solution  for  Reactions,  Chromium  Sulphate. 

164.  HC1  produces  no  precipitate. 

165.  H2S  produces  no  precipitate  in  an  acid  solution  of   a 
chromium  salt. 


82  QUALITATIVE  ANALYSIS 

1 66.  NH4OH  precipitates  bluish  green  chromium  hydroxide, 
Cr(OH)3;  soluble   with   difficulty   in    NH4OH    and  NH4C1  in 
cold  ;  reprecipitated  upon  boiling.     Soluble  in  acids. 

167.  (NH4)2S  precipitates  Cr(OH)3. 

168.  KOH  or  NaOH  precipitates  Cr(OH)3;  soluble  in  ex- 
cess, forming  green  chromite  salts,  K3CrO3  and  Na3CrO3,  repre- 
cipitated upon  boiling. 

169.  (NH^COg  or  Na2CO3  precipitates  Cr(OH)3. 

170.  The  oxides  and  the  salts  of  chromium   are  converted 
into  chromic  acid  or  chromates  by  powerful  oxidizing  agents,  — 
e.g.  by  fusion  with  sodium  carbonate  and  potassium  nitrate  on 
platinum  foil,  they  give  potassium  chromate ;  soluble  in  water. 
If    the   solution    is   acidified   with    acetic   acid,  and    lead    ace- 
tate added,   a  yellow  precipitate  of  lead  chromate  is  formed, 
PbCr<V 

171.  BaCO3  suspended  in  water,  when  shaken  with  a  cold 
neutral  or  slightly  acid  solution  of  a  chromium  salt,  precipitates 
bluish  green  chromium  hydroxide,  Cr(OH)3  ;  soluble  in   HC1, 
reprecipitated  by  167. 

CHROMIC  ACID  :  2  Solution  for  Reactions,  Potassium  Chromate. 

172.  HC1  or  other  acids  convert  yellow  potassium  chromate 
into  red  potassium  dichromate,  K2Cr2O7.     Alkalies  produce  the 
reverse  reaction  with  no  precipitation .>, 

173.  H2S  or  other  reducing  agents,  as  SO2  or  alcohol,  con- 
vert solutions  of  chromates,  to  which  HC1  has  been  added,  into 
green  solutions  of  chromium  salts. 

174.  NH4OH  produces  no  precipitate. 

1  The  oxidation  may  be  also  accomplished  by  adding  to  a  solution  of  the  salt  in  ex- 
cess of  KOH  or  NaOH,  chlorine,  bromine,  or  hydrogen  peroxide.    The  insoluble  chro- 
mium compounds  may  be  converted  to  chromates  also  by  fusion  with  NagO2  in  iron, 
nickel,  or  silver  vessels.    No  platinum  may  be  used  with  Na2O2. 

2  Reactions  for  chromic  acid  are  placed  here,  as  the  chromium  present  as  an  acid  is 
always  precipitated  as  basic  chromium  in  this  group,  and  the  reactions  for  acid  chromium 
are  very  different  from  the  reactions  of  basic  chromium. 


METAL  ANALYSIS  83 

175.  (NH^S  in   neutral  or   alkaline   solutions   precipitates 
green  chromium  hydroxide,  Cr(OH)3,  with  oxidation  of  the  sul- 
phide.    In  case  of  the  polysulphide  of  ammonia,  a  thiosulphate 
is  obtained. 

176.  (NH^COg  or  Na2CO3  produce  no  precipitate. 

177.  AgNO3  precipitates  from   neutral  solutions  red  silver 
chromate ;  soluble  in  nitric  acid  or  in  ammonia. 

178.  Lead  nitrate  or  acetate  precipitates  yellow  lead  chro- 
mate, "chrome  yellow,"  PbCrO4,  which  is  soluble  in  NaOH, 
but  with  difficulty  soluble  in  HNO3. 

179.  Hg(NO3)2  precipitates  a  dark  red  basic  mercuric  chro- 
mate, Hg2(OH)2CrO4;  soluble  with  difficulty  in  nitric  acid. 

1 80.  Bismuth  chloride  precipitates  yellowish  bismuth  chro- 
mate, Bi^CrOJgj  soluble  in  HNO3. 

181.  Cadmium  nitrate  precipitates  a  yellow  basic  cadmium 
chromate,  Cd^OH^CrC^ ;  soluble  in  HNO3. 

182.  H2O2  in  acid  solution  produces  a  deep  blue  coloration 
of  perchromic  acid,  H2Cr2O8.     If  ether  is  added  and  then  H2O2 
and  the   mixture   shaken,  the  ether  takes   up  the  perchromic 
acid,  giving  a  deep  blue  coloration  to  the  ether.1 

183.  Bead  Test?      All  compounds   of   chromium    color  the 
borax  bead  green.     See  p.  124. 

184.  BaCO3  suspended  in  water  precipitates  from  cold  neu- 
tral or  slightly  acid  solutions  yellow  barium  chromate. 

ALUMINIUM  :  Solution  for  Reactions,  Aluminium  Sulphate. 

185.  HC1  produces  no  precipitate. 

1 86.  H2S  produces  no  precipitate  from  acid  solutions. 

1  The  more  dilute  the  acid  solution  is  made,  the  more  delicate  the  test. 

2  The  platinum  wire  used  for  flame  tests  may  be  used  for  bead  tests.     Heat  the  wire 
in  the  flame  and  place  in  powdered  borax.    Heat  again  in  the  flame,  and  a  small  globule 
of  melted  borax  will  stick  to  the  wire.    Touch  while  warm  to  a  small  particle  of  the  ma- 
terial to  be  tested  and  heat  in  the  oxidizing  flame  until  the  material  has  dissolved  and 
note  the  color. 


84  QUALITATIVE  ANALYSIS 

187.  NH4OH  precipitates  white  flocculent  aluminium  hydrox- 
ide, A1(OH)3  ;  soluble  in  acids. 

1 88.  (NH4)2S  precipitates  the  same. 

189.  KOH  or  NaOH  precipitates  the  same;    soluble  in  ex- 
cess, but  reprecipitated  by  boiling  with  NH4C1.      Freshly  pre- 
cipitated A1(OH)3  is  easily  soluble  in  acids. 

190.  *(NH4)2CO3  or  Na2CO3  precipitates  A1(OH)3  or  basic 
carbonates. 

191.  Bead  Test.     No  colored  bead  is  formed. 

192.  BaCO3,  suspended  in  water,  when  shaken  with  a  cold 
neutral  or  slightly  acid  solution,  precipitates  aluminium  hydrox- 
ide, A1(OH)3. 

MANGANESE:  Solution  for  Reactions,  Manganese  Sulphate. 

193.  HC1  produces  no  precipitate. 

194.  H2S  produces  no  precipitate  from  acid  solutions. 

195.  NH4OH  in  the  absence  of  ammonium  salts  precipitates, 
incompletely,  white  Mn(OH)2;  this  oxidizes  quickly  in  the  air, 
darkening  in  color.      In  the  presence  of  ammonium  salts  this 
precipitate  is  not  formed,  but,  upon  standing,  the  solution  soon 
becomes  cloudy,  and  ultimately  all  the  manganese  is  precipi- 
tated as  dark  brown  MnO(OH)2. 

196.  (NH4)2S  precipitates  flesh-colored  manganese  sulphide, 
MnS ;  soluble  in  dilute  mineral  acids  and  acetic  acid.      It  oxi- 
dizes and  turns  brown,  upon  standing,  forming  Mn2O3,  MnSO4, 
and  sulphur. 

197.  NaOH  or  KOH  precipitates  white  Mn(OH)2;  oxidized 
quickly  by  the   air ;    darkening   in  color ;   eventually  forming 
Mn2O3. 

198.  (NH4)2CO3   or  NagCOg  precipitates  white    manganese 
carbonate   or   basic   carbonates,  which    oxidize   in   the   air   to 
Mn2O3. 

199.  Fused   with    Na2CO3   and    KNO3  upon    platinum    foil, 


METAL  ANALYSIS  85 

compounds  of  manganese  oxidize  to  a  bright  green  sodium 
manganate,  NagMnC^ ;  soluble  in  small  amount  of  water  ;  de- 
composes on  standing  in  solution  to  MnO2  and  KMnO4. 

200.  Bead  Test.    Compounds  of  manganese  color  the  borax 
bead  amethyst  in  the  oxidizing  flame ;  colorless  in  the  reducing 
flame. 

20 1.  PbO2  when  boiled  with  dilute  H2SO4  and  a  small  quan- 
tity of  manganese  salts  imparts  a  pink  or  purple  color  to  the 
solution,  due  to  the  formation  of  permanganic  acid,  HMnO4. 
The  color  is  best  noted  if  the  solution  is  allowed  to  settle  for  a 
few  minutes.     If  chlorides  are  present,  the  dry  substance  should 
first  be  treated  with   a  few  drops  of  H2SO4,  and  heated  until 
white  fumes  of  SO3  appear.     (Why  ?) 

202.  BaCO3,  suspended  in  water,  when  shaken  with  a  cold 
neutral  or  acid  solution  of  manganese  salts,  produces  no  pre- 
cipitate. 

ZINC  :    Solution  for  Reactions,  Zinc  Sulphate. 

203.  HC1  produces  no  precipitate. 

204.  H2S  precipitates  incompletely  white  zinc  sulphide,  ZnS, 
from   neutral  solutions  of  zinc  salts  of  inorganic  acids,  freely 
soluble  in  inorganic  acids ;  slightly  soluble  in  acetic  acid. 

205.  NH4OH  precipitates  white  gelatinous  zinc  hydroxide, 
Zn(OH)2  ;  soluble  in  excess  of  NH4OH  ;  reprecipitated  by  boil- 
ing.    This  precipitate  is  not  formed  in  the  presence  of  ammo- 
nium salts. 

206.  (NH^S    precipitates   white  zinc  sulphide ;   soluble  in 
dilute  inorganic  acids,  but  not  easily  in  acetic  acid. 

207.  KOH    or    NaOH    precipitates    white    gelatinous    zinc 
hydroxide,  Z^OH^;    soluble  in  excess,  forming  zincates,  as 
NajjZnOa  and  K2ZnO2. 

208.  (NH4)2CO8   or   NaaCOg  precipitates   white   basic  zinc 
carbonates ;   soluble  in  large  excess  of  reagent. 


86  QUALITATIVE  ANALYSIS 

209.  Bead  Test.     No  colored  bead  is  formed. 

210.  BaCO3,  suspended  in  water,  when  shaken  with  a  cold 
neutral  or  acid  solution  of  zinc  salts,  produces  no  precipitate. 

NICKEL  :  Solution  for  Reactions,  Nickel  Nitrate. 

211.  HC1  produces  no  precipitate. 

212.  H2S  produces  no  precipitate  from  neutral  or  acid  solu- 
tions. 

213.  NH4OH  precipitates  incompletely  light  green  nickelous 
hydroxide,  Ni(OH)2;  soluble  in  excess,  producing  a  blue  solu- 
tion.    Salts  of  ammonium  prevent  this  precipitation. 

214.  (NH4)2S    precipitates  black  nickelous    sulphide,    NiS; 
soluble  in  hot  dilute  HC1,  HNO3,  or  aqua  regia. 

215.  KOH  or  NaOH  precipitates  light  green,  Ni(OH)2 ;  in- 
soluble in  excess.     It  is  oxidized  by  boiling  with  bromine  water 
to  Ni(OH)3,  black.     If  the  precipitate  is  filtered  off  and  boiled 
with  NH4OH,  it  is  reduced  to  Ni(OH)2  with  evolution  of  nitro- 
gen.    If  NH4C1  is  present,  the  Ni(OH)2  dissolves. 

216.  (NH4)2CO3  or  Na2CO3  precipitates    light  green  basic 
carbonates  ;  soluble  in  large  excess  of  the  reagent. 

217.  KNO2  produces  no  precipitate  in  acetic  acid  solution. 
Cf.  227. 

218.  KCN    precipitates  yellowish   green  nickelous  cyanide, 
Ni(CN)2;  insoluble  in  dilute  HC1;  soluble  in  excess  of  KCN. 

219.  Bead  Test.     Salts  of  nickel  color  the  borax  bead  reddish 
brown  in  the  oxidizing  flame,  gray  in  the  reducing  flame. 

220.  BaCO3,  suspended  in  water,  when  shaken  with  a  cold 
neutral  or  acid  solution  of  nickel  salts,  produces  no  precipitate. 

COBALT  :  Solution  for  Reactions,  Cobalt  Nitrate. 

221.  HC1  produces  no  precipitate. 

222.  H2S  produces  no  precipitate  from  neutral  or  acid  solu- 
tions. 

223.  NH4OH    precipitates,  incompletely,  blue   basic  cobalt 


METAL  ANALYSIS  87 

salts  ;  soluble  in  excess  to  a  brownish  red  solution.    Ammonium 
salts  prevent  this  precipitation. 

224.  (NH4)2S  precipitates  black  cobalt  sulphide,   CoS ;    in- 
soluble in  cold  dilute  HC1;  soluble  in  HNO3  or  aqua  regia. 

225.  KOH   or  NaOH   precipitates  blue  basic    cobalt   salts, 
which  change  upon  warming  to  pink  cobalt  hydroxide,  Co(OH)2  ; 
not  soluble  in  excess  of  the  alkali.     Ammonia  or  ammonium 
salts  dissolve  the  precipitate. 

226.  Na^COg    precipitates    red    basic    cobaltous    carbonates 
which,  when  boiled,  lose  CO2,  giving  a  violet  color,  or  if  the 
reagent  is  in  excess,  a  blue  color  ;  soluble  in  ammonium  carbonate. 

227.  KNO2  precipitates,   from    solutions   rendered    strongly 
acid  with  acetic  acid,  a  yellow  crystalline  double  salt,  K3Co(NO2)6, 
but  usually  only  on  long  standing. 

228.  KCN    precipitates   brownish  white   cobaltous   cyanide, 
Co(CN)2;    soluble  in  excess  of  reagent,  and  reprecipitated  by 
HC1  or  H2SO4.     If  to  the  solution  in  excess  of   KCN  a  few 
drops  of    HC1  are  added,  and  the  solution  boiled,  potassium 
cobaltic  cyanide,  K3Co(CN)6,  is  formed,  which  is  not  reprecipi- 
tated by  sodium  hypobromite.     See  218. 

229.  Bead  Test.    Compounds  of  cobalt  color  the  borax  bead 
a  deep  blue  in  both  oxidizing  and  reducing  flame. 

230.  BaCO3,  suspended  in  water,  when  shaken  with  a  cold 
neutral  or  acid  solution  of  cobalt  salts,  produces  no  precipitate. 

Analysis,  Group  III 

Make  up  a  mixture  of  salts :  of  metals  of  this  group,  dissolve,2 
and  identify  by  the  following  scheme  of  analysis : 

A  few  drops  of  HNO3  are  added  and  the  solution  is  boiled. 
If  the  solution  turns  yellow,  iron  is  probably  present. 

1  The  salts  mentioned  on  p.  oo  are  to  be  avoided  in  making  this  solution.    In  case  the 
sample  for  analysis  contains  these  acid  radicals,  the  method  of  separation  on  p.  oo  must 
be  used. 

2  If  difficulty  is  encountered  in  making  a  solution,  consult  pp.  'oo-oo. 


88  QUALITATIVE  ANALYSIS 

(A)  Add  a  small  quantity  of  NH4C1  and  then  NH4OH  to 
alkaline  reaction  and  boil.     If  a  precipitate  forms,  one  or  more 
of  the  metals  iron,  chromium,  and  aluminium  may  be  present. 
If  no  precipitate  forms  these  are  absent,  and  subsequent  tests 
for  them  need  not  be  applied.1 

(B)  Whether  a  precipitate  is  formed  or  not  in  (A),  add  to  the 
solution  (without  filtering)  (NH4)2S  while  still  warm.     Filter  at 
once  and  wash  the  precipitate  with   water  containing  a  little 
(NH4)2S.     The  filtrate  may  contain  metals  of  Groups  IV  and 
V.     (If  the  filtrate  is  dark  brown  in  color,  the  presence  of  nickel 
is  indicated,  due  to  some  NiS  remaining  in  colloidal  solution.2 
If  this  is  the  case,  the  filtrate  should  be  acidified  with  acetic  acid 
and  boiled  for  some  time,  filtered  on  a  separate  filter,  and  the 
precipitate  tested  for  nickel  by  (D)).     The  filtrates  are  saved, 
if  analysis  of  Groups  IV  and  V  is  to  be  made. 

(C)  Remove  the  precipitate  from  the  paper  and  digest  with 
cold  dilute  HC1;3  filter,  and  wash  thoroughly. 

The  residue  may  contain  nickel  or  cobalt,  or  both  ;  test  by 
(D).  If  the  precipitate  is  light  colored,  only  sulphur  is  present 
and  may  be  discarded. 

The  filtrate  may  contain  Fe,  Al,  Cr,  Mn,  and  Zn,  and  is 
tested  by  (E). 

(D)  The  precipitate  is  tested  for  nickel  and  cobalt  with  a 
borax  bead  in  the  oxidizing  flame.     A  brown  bead  indicates 
nickel,  a  blue  bead  cobalt.     If  the  bead  is  violet  when  hot  and 
reddish  brown  when  cold,  cobalt  is  present  only  in  traces,  or  is 
absent. 

If  cobalt  is  present,  nickel  may  also  be  present,  but  the  test 
for  nickel  is  obscured  by  the  deep  color  of  the  cobalt  bead. 

1  If  desired,  this  precipitate  may  be  filtered  and  examined  as  in  (G),  (H),  and  (I).    The 
filtrate  is  then  treated  with  (NH4)2S  and  treated  as  in  (C),  (D),  and  (L). 

2  If  the  precipitation  has  been  performed  properly,  no  NiS  should  remain  in  solution. 
8  See  Appendix  for  making  dilute  HC1. 


METAL  ANALYSIS  89 

If  cobalt  is  present,  proceed  as  follows  :  To  the  precipitate 
add  a  small  amount  of  HC1  and  a  few  drops  of  HNO3,  and 
evaporate  nearly  to  dryness.  Add  about  5  c.c.  of  water  and 
filter ;  discard  the  residue  on  the  filter,  which  is  sulphur. 
Evaporate  to  a  small  bulk,  if  necessary.  Add  NaOH,  drop  by 
drop,  until  a  permanent  precipitate  is  formed ;  this  is  dissolved 
in  a  slight  excess  of  KCN  1  and  boiled.  A  volume  of  bromine 
water  equal  to  the  whole  solution  is  now  added,  and  an  excess 
of  NaOH.  If  a  precipitate  is  formed,  it  is  filtered,  washed,  and 
tested  for  nickel  with  the  borax  bead. 

Concentrate  the  filtrate  to  a  small  volume.  Test  solution  with 
borax  bead  for  Co. 

(E)  The  filtrate  from  the  sulphides  of  nickel  and  cobalt  is 
boiled  with  a  few  drops  of  HNO3  to  oxidize  the  iron  and  is  fil- 
tered if  necessary ;  any  precipitate  which  may  be  formed  is  dis- 
carded since  it  is  sulphur.     (Iron  is  tested  for  in  the  filtrate  by 
taking  a  few  drops  and  adding  a  solution  of  ammonium  sulpho- 
cyanate.     A  deep  red  color  indicates  iron.     A  light  red  color 
shows  that  iron  is  present  only  in  traces.)     Add  Na2CO3  to  the 
solution  with  stirring  until  a  permanent  precipitate  just  com- 
mences to  form.     Dissolve  with  a  drop  of  dilute  HC1.     Add 
about  two  drops  of  acetic  acid  and  an  excess  of  sodium  acetate,2 
and  boil  until,  on  standing,  the  solution  above  the  precipitate  is 
clear.3     Filter  hot.     The  precipitate  contains  the  hydroxides  of 
any  iron,  chromium,  and  aluminium  which  may  be  present ;  to 
be  determined  by  (F).     The  filtrate  contains  any  manganese  or 
zinc  which  may  be  present ;  to  be  determined  by  (K). 

(F)  The  precipitate  is  dissolved  with  dilute  HC1,  treated  with 

1  A  large  excess  of  KCN  tends  to  prevent  precipitation  of  nickel,  should  it  be  present. 

2  Five  cubic  centimeters  may  be  added  and,  if  a  precipitate  is  formed  upon  boiling, 
more  may  be  added  so  long  as  further  addition  produces  a  turbidity  of  the  clear  liquid. 
If  the  first  addition  produces  no  precipitate  upon  boiling,  further  addition  is  unnecessary. 

3  The  hydrolysis  of  the  acetates  of  Al,  Cr,  and  Fe  is  analogous  to  that  of  the  carbonates 
and  sulphides  of  the  same  metals.  Cf.  reactions  157,  167,  169,  175,  and  176.  See  also  p.  36. 


90  QUALITATIVE  ANALYSIS 

an  excess  of  NaOH,  boiled,  and  filtered.  The  precipitate  may 
be  a  mixture  of  iron  and  chromium  hydroxides. 

The  filtrate  may  contain  aluminium. 

(G)  Fuse  on  a  platinum  foil  with  Na2CO3  and  KNO3  ;  a  yellow 
color  indicates  chromium.  Boil  with  water,  filter,  and  to  the 
filtrate  add  a  few  drops  of  acetic  acid,  to  acid  reaction,  and  lead 
acetate.  Yellow  lead  chromate  is  formed  if  chromium  is  present. 

Alternative  test  for  chromium  :  fuse  the  precipitate  from  (F) 
with  sodium  peroxide  in  a  nickel,  iron,  or  silver  vessel,  dissolve 
in  water,  acidify  with  acetic  acid,  and  test  as  before.  (See  also 
reaction  170  and  footnote.) 

(H)  The  residue  from  (G)  insoluble  in  water  is  dissolved  in 
hydrochloric  acid,  and  tested  for  iron  with  ammonium  sulpho- 
cyanate. 

(I)  The  nitrate  from  the  NaOH  solution  (F)  is  acidified  with 
HC1,  an  excess  of  NH4OH  added,  heated,  and  allowed  to  stand 
a  few  minutes.  A  white  flocculent  precipitate  shows  presence 
of  aluminium.1 

(K)  The  filtrate  (E)  may  contain  zinc  and  manganese.  Acidify 
with  acetic  acid  and  conduct  H2S  into  the  boiling  solution.  White 
zinc  sulphide  is  precipitated.  Filter,  test  the  precipitate  for  zinc 
with  dilute  HC1.  If  soluble,  the  presence  of  zinc  is  shown.  The 
filtrate  from  the  ZnS  precipitate  is  evaporated  to  dryness,  and 
divided  into  two  portions.  To  one  add  two  drops  of  cone.  H2SO4 
and  heat  until  white  fumes  of  SO3  are  given  off.  Cool,  add  dilute 
H2SO4,  and  transfer  to  a  test  tube,  and  fill  half  full  with  dilute 
H2SO4.  Add  PbO2,  and  heat  to  boiling,  and  allow  to  stand 
until  settled.  A  purple  or  red  solution  indicates  manganese. 

The  other  portion  is  mixed  with  NagCOg  and  KNO3  and  heated 
on  a  platinum  foil.  A  green  color  indicates  manganese. 

1  A  flocculent  precipitate  may  be  due  to  silica  derived  from  impure  sodium  hydroxide. 
The  test  may  be  confirmed  by  filtration  and  solution  of  the  precipitate  in  dilute  hydro- 
chloric acid. 


METAL  ANALYSIS  91 

After  having  completed  the  analysis  of  the  mixture  proceed  to 
the  analysis  of  unknowns  Nos.  7  and  8.  These  being  reported, 
prepare  and  record  the  following : 

EXERCISES 

I .  What  reactions  are  com  mon  to  all  the  members  of  this  group  ? 

II.  Devise  another  method  for  the  separation  of  nickel  and 
cobalt. 

III.  Give  the  name  and  composition  of  at  least  one.  ore  of  each 
of  the  members  of  this  group. 

IV.  For  what  purposes  are  the  following  substances  useful, 
aside   from  analytical  work :    ferric  chloride,   potassium   ferro- 
cyanide,  Prussian  blue,  potassium  chromate,    potassium   alum, 
potassium  permanganate,  pyrolusite,  zinc  oxide,  cobalt  nitrate  ? 

V.  What  is   steel  ?     Cast  iron  ?     Wrought  iron  ?     Spiegel- 
eisen  ? 

VI.  What  are  the  chief  uses  of  each  of  the  metals  of  this 
group  ? 

VII.  Why  cannot  you  get  a  test  for  iron  in  K4Fe(CN)6  by  the 
methods  given  for  this  group  ? 

VIII.  Why  are  the  carbonates,  sulphides,  and  acetates  of  ferric 
iron,  aluminium,  and  chromium  decomposed  by  water  ? 

IX.  Why  does  ammonium  hydroxide  fail  to  precipitate  man- 
ganese when  ammonium  salts  are  present? 

X.  Why  must  organic  compounds  be  removed  before  this 
group  can  be  analyzed. 

Group  IV 

The  Ammonium  Carbonate  Group 
Barittm,  Strontium,  and  Calcium 

GENERAL  STATEMENT.  The  carbonates  of  this  group 
are  insoluble  in  water  and  in  solutions  of  ammonium  salts. 
Magnesium  carbonate  is  insoluble  in  water,  but  soluble  in  a 


92  QUALITATIVE  ANALYSIS 

solution  of  ammonium  chloride,  and  is  placed  in  Group  V  (the 
soluble  group). 

BARIUM  :  Solution  for  Reactions,  Barium  Chloride. 

231.  HC1,  H2S,  NH4OH,  and  (NH4)2S  produce  no  precipitate. 

232.  KOH  or  NaOH  precipitates  from  concentrated  solutions 
white  voluminous  barium  hydroxide,  Ba(OH)2;  insoluble  in  ex- 
cess of  reagent,  but  soluble  in  water. 

233.  (NH4)2CO3  or  Na2CO3  precipitates  barium  carbonate, 
BaCO3 ;  soluble  in  acids,  even  acetic  acid. 

234.  Na2HPO4  precipitates  acid  barium  phosphate,  BaHPO4, 
easily  soluble  in  dilute  HC1  or  HNO3. 

235.  K2CrO4  precipitates  yellow  barium  chromate,  BaCrO4 ; 
soluble  in  HC1;  insoluble  in  acetic  acid. 

236.  (NH4)2C2O4  precipitates  white  barium  oxalate,  BaC2O4 ; 
soluble  in  acetic  acid. 

237.  H2SiF6  precipitates  white  barium  fluosilicate,  BaSiF6 ; 
insoluble  in  alcohol. 

238.  Sulphuric  acid  or  soluble   sulphates  precipitate  white 
barium  sulphate,  BaSO4;  insoluble  in  acids. 

239.  CaSO4  or  SrSO4  precipitates  BaSO4. 

240.  Flame  Test.1     Placed  upon  a  platinum  wire  and  intro- 
duced into  a  Bunsen  flame,  a  yellowish  green  coloration  is  pro- 
duced. 

STRONTIUM  :  Solution  for  Reactions,  Strontium  Chloride. 

241.  HC1,  H2S,  NH4OH,  (NH4)2S  produce  no  precipitate. 

242.  KOH  or  NaOH  precipitates  from  concentrated  solutions 
white  voluminous  Sr(OH)2,  resembling  Ba(OH)2  (q.v.\  but  less 
soluble  in  water. 

1  For  the  flame  test  the  platinum  wire  should  be  cleaned  by  dipping  in  a  solution  of  HC1 
and  heating  repeatedly  until  no  color  is  given  to  the  flame.  The  wire  is  then  dipped  into 
the  salt  and  introduced  into  the  flame.  If  the  color  does  not  show  immediately,  dip  into 
HC1  and  try  again.  Some  salts  do  not  give  a  good  flame  test  until  after  being  treated  with 
HC1  on  the  platinum  wire.  (Why  ?) 


METAL  ANALYSIS  93 

243.  (NH^COg  or  Na2CO3  precipitates  white  strontium  car- 
bonate, SrCOg,  resembling  BaCO3 ;  easily  soluble  in  acids,  even 
acetic. 

244.  Na2HPO4  precipitates  white  acid  strontium  phosphate, 
SrHPO4;  easily  soluble  in  dilute  HNO3  or  HC1. 

245.  K2CrO4  precipitates  only  in  concentrated  solutions  yellow 
strontium  chromate.     The  presence  of  acetic  acid  prevents  the 
precipitation. 

246.  ( N H4)2C2O4  precipitates  white  strontium  oxalate,  SrC2O4 ; 
only  slightly  soluble  in  acetic  acid. 

247.  H2SiF6  produces  no  precipitate  even  with  the  addition  of 
alcohol. 

248.  H2SO4  precipitates  white  strontium  sulphate,   SrSO4; 
slightly  soluble  in  water  and  hence  does  not  appear  at  once  in 
dilute  solutions.     More  soluble  than  barium  sulphate. 

249.  CaSO4  precipitates  strontium  sulphate. 

250.  Flame  Test.     Placed  upon  a  platinum  wire  and  intro- 
duced into  a  Bunsen  flame,  a  brilliant  crimson  color  is  produced. 

CALCIUM  :  Solution  for  Reactions,  Calcium  Chloride. 

251.  HC1,  H2S,  NH4OH,  (NH^S  produce  no  precipitate. 

252.  KOH  or   NaOH  precipitates  from  sufficiently  concen- 
trated solutions  CaCOH^;  less  soluble  than  S^OH^  (q.v.). 

253.  (NH^COg  or  Na^COg  precipitates  white  calcium  car- 
bonate, CaCO3;  slightly  soluble  in  an  excess  of  concentrated 
solution  of  Na^Og ;  easily  soluble  in  acids,  even  acetic. 

254.  Na2HPO4  precipitates   white   acid   calcium   phosphate, 
CaHPO4;  easily  soluble  in  dilute  HC1  or  HNO3. 

255.  K2CrO4  produces  no  precipitate. 

256.  (NH4)2C2O4  precipitates  white  calcium  oxalate,  CaC2O4; 
insoluble  in  acetic  acid;  soluble  in  HC1  or  HNO3. 

257.  H2SiF6  produces  no  precipitate,  even  with  addition  of 
alcohol. 


94  QUALITATIVE  ANALYSIS 

258.  H2SO4  precipitates  from  strong  solutions  calcium  sul- 
phate, CaSO4 ;  more  soluble  in  water  than  BaSO4  or  SrSO4. 

259.  Flame  Test.     Brought  upon  a  platinum  wire  and  intro- 
duced into  a  flame,  compounds  of  calcium  produce  a  brick-red 
flame.     Compare  with  strontium. 

Analysis,  Group  IV 

Make  a  mixture  of  salts  of  the  metals  of  this  group.  Dissolve 
in  water  with  the  addition  of  a  little  HC1,  if  necessary,  and  iden- 
tify by  the  following  scheme  for  analysis  : 

(A)  Ammonium  chloride,  ammonium  hydroxide,  and  ammo- 
nium carbonate  are  added  and  the  solution  warmed  for  about  ten 
minutes.     Filter  and  wash  with  hot  water.     The  filtrate  may  con- 
tain members  of  Group  V.     A  small  portion  of  the  precipitate 
should  be  reserved  for  flame  tests.     Dissolve  the  remainder  in 
dilute  acetic  acid. 

(B)  To  a  small  portion  of  the  acetic  acid  solution  add  K2Cr2O7. 
If  no  precipitate  forms,  proceed  at  once  with  the  remainder  of  the 
solution  to  (E)  and  (F).     If  a  precipitate  forms,  treat  the  whole 
solution  with  dichromate  and  filter.     Wash  the  precipitate  and 
test  by  (C). 

(C)  Dissolve  the  precipitate  by  warming  in  dilute  HC1  and 
add  dilute  H2SO4.     A  precipitate  indicates  barium.     The  pre- 
cipitate will  appear  yellow  until  filtered  and  washed. 

(D)  The  filtrate  from  (B)  is  made  alkaline  with  NH4OH  and 
reprecipitated  with  (NH4)2CO3.     If  no  precipitate  forms,  Ca  and 
Sr  are  both  absent.      If  a  precipitate  forms,  filter  and  wash 
until  white.     (The  precipitate  may  be  tested  for  strontium  and 
calcium  by  the  flame  test.     See  250  and  259.)     Dissolve  the  pre- 
cipitate in  acetic  acid  and  divide  into  two  portions  and  analyze 
by  (E)  and  (F). 

(E)  To  one  portion  of  the  solution  from  (D)  add  a  saturated 
solution    of    CaSO4.      Stir  vigorously  and   allow  to   stand  for 


METAL  ANALYSIS  95 

some  time.  A  precipitate  indicates  Sr.  Confirm  by  flame 
test. 

(F)  To  the  other  portion  of  the  solution,  add  a  solution  of 
NagSC^  until  precipitation  is  complete.  Filter,  and  to  the  ni- 
trate add  a  solution  of  (NH4^C2O4.  A  precipitate  indicates 
the  presence  of  calcium.  Confirm  by  flame  test.  If  no  stron- 
tium is  present,  the  treatment  with  NagSC^  may  be  omitted. 

After  having  completed  the  analysis  of  a  known  mixture,  pro- 
ceed to  the  analysis  of  unknowns  Nos.  9  and  10.  These  being 
reported,  prepare  and  record  the  following : 

EXERCISES 

I.  What  reactions  are  common  to  all  the  members  of  this 
group  ? 

II.  Prepare  a  list  of  the  more  important  minerals  containing 
the  metals  of  this  group. 

III.  How  is  lime  made  ?    What  is  limewater  ?     Baryta  water? 
Bleaching  powder  ?     Superphosphate  of  lime  ?     Plaster  of  Paris  ? 

IV.  How  are  the  peroxides  of   barium,  calcium,  strontium, 
and  magnesium  prepared  ?     What  are  the  uses  of  each  ? 

V.  What  is  the  composition  of  window  glass  ?     Hard  glass  ? 
Jena  glass  ? 

VI.  What  are  the  commercial  uses  of  heavy  spar  ? 

VII.  What  is  the  composition  of  plaster? 

VIII.  What  use  is  made  of  the  polysulphide  of  calcium  and 
how  is  it  prepared  ? 

Group  V 

The  Soluble  Group 
Sodium,  Potassium,  Ammonium  (the  Alkalies},  and  Magnesium 

GENERAL  STATEMENT.  As  there  is  no  reagent  which 
will  precipitate  all  the  members  of  this  group,  it  is  known  as 
the  soluble  group. 


96  QUALITATIVE  ANALYSIS 

SODIUM:  Solution  for  Reactions,  Sodium  Chloride. 

260.  K2H2Sb2O7  precipitates  in  fairly  concentrated  neutral  or 
weakly  alkaline  solution,  Na2H2Sb2O7.     Rubbing  with  a  glass 
rod  facilitates  precipitation.     In  acid  solution,  H2Sb2O7  is  pre- 
cipitated. 

261.  K3A1F6   produces   a   flocculent  precipitate,   Na3AlF6. 1 
The  same  precipitate  is  also  produced  by  magnesium  salts,  but 
not  by  potassium  and  ammonium  salts. 

262.  Flame  Test.     When  a  sodium  salt  is  brought  upon  a 
platinum  wire  and  introduced  into  a  Bunsen  flame,  an  intense 
yellow   coloration  is  produced,  which  is  not  visible  through  a 
blue  glass.2 

POTASSIUM  :  Solution  for  Reactions,  Potassium  Chloride. 

263.  PtCl4  precipitates,  except  in  very  dilute  solutions,  yellow 
crystalline  potassium  chlorplatinate,  K2PtCl6 ;  insoluble  in  alco- 
hol. 

264.  NaHC4H4O6    or    H2C4H4O6    precipitates    white    acid 
potassium  tartrate,  KHC4H4O6.     Agitating,  or  scratching  the 
walls  of  the  containing  vessel  with  a  glass  rod,  facilitates  the 
precipitation. 

265.  Na3Co(NO2)6   precipitates    from    neutral    concentrated 
solutions  potassio-cobaltic  nitrite,  K3Co(NO2)6,  in  the  form  of  a 
yellow  powder.     Precipitation  takes  place  slowly  from   dilute 
solutions,  but  is  hastened  by  warming  gently.     If  the  solution 
is  alkaline,  it  should  be  made  slightly  acid   with  acetic  acid 
before  applying  the  test.     If  acid,  it  should  be  made  slightly 
alkaline  with  sodium  carbonate  and  then  slightly  acidified  with 
acetic  acid. 

266.  Flame  Test.     Placed  upon  platinum  wire  and  heated  in 
a  Bunsen  flame,  compounds  of  potassium  color  the  flame  violet. 

1  The  ratio  of  Na  to  Al  is  i.i :  i. 

2  The  color  is  not  visible,  but  the  student  should  note  the  distortion  of  the  flame  by  the 
wire  and  compare  this  with  the  appearance  of  the  flame  in  reaction  266. 


METAL  ANALYSIS  97 

A  small  amount  of  sodium  salts  obscures  this  test,  but  if  the 
flame  is  observed  through  a  cobalt  glass,  the  sodium  rays  are 
cut  off  and  the  potassium  appears  reddish  violet. 

AMMONIUM  :    Solution  for  Reactions,  Ammonium  Chloride. 

267.  PtCl4  precipitates  in  concentrated  solutions  yellow  crys- 
talline   ammonium    chlorplatinate,    (NH^PtClg,    insoluble    in 
alcohol. 

268.  NaHC4H4O6  or  H2C4H4O6  precipitates  white  acid  am- 
monium tartrate,  NH4HC4H4O6,  from  quite  concentrated  solu- 
tions upon  standing.     The  precipitation  is  hastened  by  shaking, 
or  scratching  the  walls  of  the  containing  vessel  with  a  glass  rod. 

269.  Na3Co(NO2)6    precipitates    from    acetic    acid    solution 
ammonium  cobaltic  nitrite,  (NH4)3Co(NO2)6,  in  the  form  of  a 
yellow  powder. 

270.  NaOH   or   KOH  in  excess  liberates   ammonia,    NH3. 
Recognized  by  (a)  odor,  (b)  action  of  the  gas  on  moist  litmus, 
changing  red  litmus  to  blue,  (c)  fumes  produced  when  a  glass 
rod  moistened  with  HC1  is    held  in  escaping    gas.     Action  is 
hastened  by  warming. 

271.  Ca(OH)2  liberates  ammonia  gas.     If  solid  calcium  oxide, 
CaO,  and  an  ammonium  salt  are  mixed  in  a  beaker,  the  test  for 
free  ammonia  may  be  made  by  placing  over  it  a  watch  glass,  upon 
the  under  side  of  which  is  placed  a  piece  of  moist  red  litmus. 

272.  Nessler's  Reagent1  is  a  most  delicate  test2  for  ammonia. 
It  gives  a  distinct  yellow  coloration  to  solution  containing  minute 
quantities  of  ammonium  compounds.     If  present  in  large  quan- 
tities, a  brown  precipitate  is  produced. 

273.  All  ammonium  compounds  are  volatilized  by  heat. 

274.  Flame  Test.     Placed  upon  a  platinum  wire  and  intro- 
duced into  a  Bunsen  flame,  ammonium  compounds  give  a  reddish 

*  See  Appendix. 

2  The  test  is  so  very  delicate  as  to  be  unsuited  to  ordinary  qualitative  operations. 


98  QUALITATIVE  ANALYSIS 

violet  flame  when  viewed  through  a  cobalt  glass.     Compare  with 
the  potassium  flame. 

MAGNESIUM  :  Solution  for  Reactions,  Magnesium  Sulphate. 

275.  PtCl4,  Na3Co(NO2)6,  H2C4H4O6,  NaHC4H4O6  and  Nes- 
sler's  solution  give  no  tests  with  magnesium  salts. 

276.  NH4OH     precipitates     white     magnesium     hydroxide, 
Mg(OH)2  ;  soluble  in  presence  of  ammonium  salts. 

277.  (NH4)2CO3    precipitates  white    magnesium    carbonate, 
MgCO3 ;  from  concentrated  solution  ;  soluble  in  presence  of  am- 
monium salts. 

278.  KOH  or  NaOH  precipitates  white  magnesium  hydrox- 
ide, Mg(OH)2;  almost  insoluble  in  water. 

279.  Na2HPO4  in    presence   of    NH4OH   and   NH4C1   pre- 
cipitates white  crystalline    magnesium    ammonium    phosphate, 
MgNH4PO4.     The   precipitation   is   slow   in   dilute   solutions ; 
hastened  by  warming  and  agitation. 

280.  H2SO4,  H2SiF6,  and  (NH4)2C2O4  produce  no  precipitate. 

Analysis,  Group  V 

Make  a  mixture  of  Na,  K,  NH4,  and  Mg  salts.  Dissolve  a 
portion  of  the  mixture  and  identify  by  the  following  scheme  for 
analysis  : 

(A)  To  a  portion  of  the  solution,  add  a  small  excess  of 
NH4OH,  and  some  NH4C1  in  case  it  is  not  already  known  to 
be  present,  and  Na2HPO4.  Thoroughly  agitate,  or  scratch  with 
a  glass  rod,  if  a  precipitate  does  not  form  at  once,  and  allow  to 
stand  for  some  hours.  A  white  crystalline  precipitate  indicates 
magnesium.1 

1  In  case  an  unknown  material  is  being  analyzed,  it  may  contain  other  cations  besides 
those  of  this  group.  If  this  be  true,  they  may  be  removed  by  treatment  with  ammonium 
carbonate,  ammonium  hydroxide,  and  ammonium  chloride,  and  the  filtrate  used  as  directed 
above. 


METAL  ANALYSIS  99 

(B)  A  second  portion  of  the  solution  is  evaporated  to  dryness 
and  heated  to  a  low  red  heat  in  a  porcelain  dish  until  white  fumes 
are  no  longer  observed.     Care  must  be  taken  to  heat  all  portions 
of  the  dish.     When  cool,  dissolve  in  as  small  an  amount  of  water 
as  possible,  add  a  few  drops  of  HC1,  and  filter  if  necessary.     Test 
the  filtrate  for  potassium  with  Na3Co(NO2)6  or  PtCl4,  and  with 
the  flame  test,  using  cobalt  glass. 

(C)  A  third  portion  of  the  solution  may  be  treated  with  K3A1F6. 
A  precipitate  indicates  sodium.    See  261.    The  flame  test  may  be 
used  on  some  of  the  solid  material  obtained  in  (B)  or  on  the  origi- 
nal dry  substance.     A  persistent  yellow  flame  indicates  sodium.1 

(D)  A  portion  of  the  original  material  is  treated  with  lime  or 
a  solution  of  KOH  and  the  ammonia  sought  by  (a)  odor,(£)  litmus, 
(c)  action  on  a  rod  moistened  with  HC1.     See  270  and  271. 

After  completing  the  analysis  of  a  known  mixture,  the  stu- 
dent will  proceed  to  the  analysis  of  samples  Nos.  1 1  and  12,  after 
which  the  following  exercises  will  be  prepared  and  recorded  : 

EXERCISES 

I.  Devise  another  method  for  the  analysis  of  this  group. 

II.  What  are  the  more  important  salts  of  ammonium?   of 
magnesium  ?   of  potassium  ?    of  sodium  ? 

III.  What  are  the  more  important  compounds  of  these  metals 
found  in  nature  ? 

IV.  Describe  the  Solvay  process  for  the  manufacture  of  soda ; 
the  Leblanc  process. 

V.  What  is  Glauber's  salt  ?     Epsom  salt  ?    Saleratus  ?    Chili 
saltpetre  ?     Sal  ammoniac  ?     Washing  soda  ?     Gunpowder  ? 

VI.  Describe  briefly  the  usual  impurities  found  in  common 
salt,  and  its  purification. 

1  The  flame  test  for  sodium  is  so  very  delicate,  and  the  presence  of  sodium  in  small 
quantities  is  so  frequent,  that  in  reporting  unknowns,  considerable  discrimination  must  be 
used. 


100  QUALITATIVE  ANALYSIS 

VII.  What  substances  are  present  in  a  solution  of  ammonium 
hydroxide?     See  Smith's  General  Chemistry,  p.  338  and  p.  565. 

VIII.  Write  a  detailed  explanation  of  the  effect  of  ammonium 
salts  on  the  precipitation  of  magnesium  by  ammonium  hydroxide. 
Smith's  General  Chemistry,  p.  644. 

IX.  What  substances  are  present  in  a  solution  of  pure  sodium 
chloride  ? 


PART   III 

ACID  ANALYSIS 

GENERAL  CONSIDERATIONS.  It  is  usual  to  make 
use  of  the  negative  ions  of  acids  or  of  salts  in  the  detection  of 
the  metalloids  and  non-metals.  If  the  non-metal  does  not 
already  exist  in  a  negative  ion,  it  is  converted  into  such  and  its 
identity  determined  by  the  characteristic  appearance  or  behavior. 
Acid  analysis  is,  then,  the  detection  of  the  non-metals.  Also 
negative  groups  may  at  times  involve  elements  which  also  reveal 
themselves  as  cations  during  metal  analysis,  e.g.  arsenic,  man- 
ganese, chromium,  etc. 

Just  as  in  the  case  of  the  analysis  for  cations,  it  is  essential 
to  acid  analysis  that  the  substance  be  brought  into  solution  in 
order  to  facilitate  identification.  The  appearance  of  this  solu- 
tion, or  the  absence  of  certain  metals,  may  at  times  render  the 
search  for  certain  anions  unnecessary,  e.g.  a  colorless  solution 
is  sufficient  evidence  of  the  absence  of  chromates  or  of  per- 
manganates, while  the  failure  to  find  arsenic  in  the  course  of 
the  metal  analysis  renders  search  for  arsenic  and  arsenous  acids 
futile.  Again  the  presence  of  lead  or  of  barium  in  a  soluble 
substance  precludes  the  possibility  of  the  presence  of  sulphuric 
acid.  In  order  to  assist  in  forming  such  conclusions,  as  well  as 
for  other  purposes,  the  following  rules  of  solubility  will  be  found 
useful. 

NOTE.    See  also  Table  of  Solubilities,  Appendix. 

(i)  All  sodium,  potassium,  and  ammonium  salts  are  soluble 
in  water,  except  the  chlorplatinates  and  acid  tartrates  of  ammo- 

101 


102  QUALITATIVE  ANALYSIS 

nium  and  potassium,  the  silicofluoride  of  sodium  and  potassium, 
and  the  aluminofluoride  of  sodium. 

(2)  All  chlorates,  nitrates,  and  acetates  are  soluble  in  water 
(basic  salts  excepted). 

(3)  All  carbonates,  phosphates,  borates,  oxalates,  and  arse- 
nates,  except  those  of  the  alkalies,  are  insoluble  in  water,  but 
are  soluble  in  dilute  acids. 

(4)  All  chlorides,  bromides,  and  iodides  are  soluble  in  water, 
except  those  of  silver,  lead,  and  mercury  (ous),  and    mercuric 
iodide. 

(5)  All  sulphates  are  soluble  except  those  of  barium,  stron- 
tium, and  lead.     (Calcium  sulphate  and  silver  sulphate  are  but 
slightly  soluble.) 

(6)  All  hydroxides  are  insoluble  except  those  of  the  alkalies 
and  alkaline  earths. 

It  may  further  be  remembered  that  the  nature  of  the  sub- 
stance limits  the  search  for  acids,  e.g.  an  alloy  can  contain  no 
acids  and  but  a  limited  number  of  metals.  An  alloy  may, 
however,  contain  non-metallic  elements  such  as  C,  S,  Si,  P,  etc. 
A  mineral  insoluble  in  water  contains  no  organic  acids  or 
cyanogen  compounds.  It  likewise  contains  no  nitrates,  chlo- 
rates, etc. 

In  the  qualitative  analysis  for  acids  there  is  no  absolute  sepa- 
ration into  groups,  and  subsequent  separation  of  each  group, 
either  possible  or  desirable.  Nevertheless  the  acids  may  be 
advantageously  grouped  and  group  reagents  suggested,  so  that 
the  presence  of  representatives  of  each  group  may  be  affirmed, 
or  denied,  by  single  tests.  Other  than  in  this  respect  the  iden- 
tification of  the  acids  is  by  means  of  individual  tests.  In  these 
notes  only  the  acids  more  frequently  encountered  will  be  dis- 
cussed. The  rare  elements  which  form  acids  are  briefly  dis- 
cussed in  Part  IV. 

The  following  grouping  is  suggested  : 


ACID  ANALYSIS  103 

Group  I 

Acids  the  anions  of  which  are  precipitated  from  salts  by 
dilute  nitric  acid  or  are  decomposed  by  nitric  acid :  H2SiO3, 
H2S.HNO2,  H2S2O3,  H2SO3.H2CO3,  HC1O. 

Group  II 

Acids  the  anions  of  which  are  precipitated  in  dilute  nitric  acid 
solution  by  silver  nitrate  : 

HC1,  HBr,  HI,  HCN,  H3Fe(CN)6.  H4Fe(CN)6. 

Group  III 

Acids  the  anions  of  which  are  precipitated  by  barium  chloride 
from  neutral  solutions:  H2SO4 .  H2CrO4,  H3PO4,  H3BO3, 
H2C204,  H2C406,  H3As03,  H3AsO4,  HF,  H4SiO4. 

Group  IV 

Acids  the  anions  of  which  are  not  precipitated  by  silver 
nitrate  or  by  barium  chloride:  HNO3,  HC1O3,  HC2H3O2, 
HMnO4. 

Group  V 

Organic  acids.  Almost  all  of  the  organic  acids  give  a  "  burnt " 
odor  on  heating  to  redness  and  give  a  black,  charred  residue. 
None  of  these  will  be  considered  in  this  connection  save  tartaric 
acid,  oxalic  acid,  and  acetic  acid.  These  have  been  included  in 
groups  three  and  four.  Oxalic  acid  neither  chars  nor  gives  a 
burnt  odor  on  heating. 

Preliminary  Reactions 

Students  are  advised  to  perform  all  preliminary  reactions 
with  known  solutions  and  to  write  equations  representing  the 
same  before  proceeding  to  the  analysis  of  any  unknown  sub- 
stance. For  this  purpose,  it  is  best  to  use  solutions  of  the  alkali 
salts  of  the  acids. 


104  QUALITATIVE  ANALYSIS 

Group  I 
H2SiO3,  H2S,  HNO2,  H2S2O3,  H2SO3,  H2CO3,  HC1O. 

SILICIC  ACID 

Silicic  acid  is  precipitated  from  silicates  soluble  in  water, 
under  most  circumstances,  by  acids  as  a  jelly  like  transparent 
mass  which  always  leaves  an  insoluble  residue  of  silicon  dioxide 
when  evaporated  to  dryness  and  heated  to  no  degrees  centi- 
grade. This  residue  may  be  tested  for  silicon  by  reaction  2 
or  3,  p.  115.  For  the  behavior  of  insoluble  silicates,  see 
Group  III. 

HYDROSULPHURIC  ACID 

1.  Dilute  hydrochloric  acid  evolves  hydrogen  sulphide  from 
most  sulphides  and  will  darken  a  bit  of  filter  paper  moistened 
with  an  alkaline  solution  of  lead  acetate.     (See  Appendix.)     If 
the  sulphide  is  insoluble,  it  will  give  the  above  test  when  the 
powdered  solid  is  mixed  with  zinc  dust  and  then  treated  with 
acid. 

2.  Silver  nitrate  precipitates  silver  sulphide  from  hydrogen 
sulphide  or  soluble  sulphides. 

3.  Sodium  nitroprusside,  in  alkaline  solutions,   gives  a  red- 
dish violet  color  with  negative  sulphur  ions. 

4.  Concentrated    sulphuric    acid    decomposes    all    sulphides 
with  liberation  of  sulphur  dioxide  and  sulphur. 

5.  Antimony  salts  precipitate  red  antimony  trisulphide. 

SULPHUROUS  AND  THIOSULPHURIC  ACIDS 

I.  Dilute  sulphuric  acid  or  hydrochloric  acid  evolves  sulphur 
dioxide  from  all  sulphites  in  the  cold.  Thiosulphates  when 
acidified  also  liberate  sulphur  dioxide  as  well  as  a  milky  pre- 
cipitate of  sulphur. 


ACID  ANALYSIS  105 

2.  Acid  solutions  of  chromates  and  permanganates  are  de- 
composed by  sulphites  and  sulphur  dioxide ;  thiosulphates  also 
liberate  sulphur. 

3.  Iodine  solutions  are  decolorized  by  sulphurous  acid. 

4.  Strontium  chloride  produces  a  precipitate  in  neutral  solu- 
tions of  sulphites,  but  not  in  thiosulphates  unless  very  concen- 
trated.    (Strontium  thiosulphate  is  soluble  in  3.7  parts  water.) 

A  small  precipitate  is  nearly  always  produced  in  solutions  of 
commercial  salts,  since  they  nearly  always  contain  some  sul- 
phate. 

5.  Ferric  chloride  produces  a  dark  violet  evanescent  colora- 
tion with  thiosulphates  but  not  with  sulphites. 

6.  On  boiling  with  bromine  water  sulphites  and  thiosulphates 
are  converted  into  sulphates,  which  are  precipitated  from  acid 
solution  by  barium  chloride. 

CARBONIC  ACID 

1.  Dilute  acids  decompose  all  carbonates,  most  of  them  in 
the  cold.     When  the  gas  evolved  is  passed  into  limewater  or 
baryta  water,  a  deposit  of  carbonate  is  obtained. 

2.  Concentrated  acids  behave  in  like  manner,  but  more  vigor- 
ously.    (See  also  organic  acids.) 

NITROUS  ACID 

1.  Ferrous   sulphate  and  concentrated  sulphuric  acid   give 
nitrosyl  ferrous  sulphate,  as  in  case  of  nitrates  (q.v.). 

2.  Dilute  acids  react  on  nitrites  with  the  production  of  brown 
vapors. 

3.  Indigo  and  permanganates  are  decolorized  by  nitrites  in 
acid  solutions. 

4.  Potassium  iodide  treated  with  nitrites  in  acid  solution  lib- 
erates iodine ;  detected  by  shaking  with  chloroform  or  carbon 
disulphide. 


106  QUALITATIVE  ANALYSIS 

HYPOCHLOROUS  ACID 

1.  All  acids,  even  carbonic,  liberate  chlorine  from  hypochlo- 
rites.     All  hypochlorites  are  soluble  in  water. 

2.  lodo-starch    paper   is   colored   blue   by   hypochlorites   in 
weakly  alkaline  solutions  by  reason  of  the  liberation  of  iodine. 

Analysis  of  Group  I 

Analysis  of  Nos.  i  and  2  of  a  set  of  unknowns  is  to  be  un- 
dertaken, and  the  outline  given  below  should  be  followed. 

Treat  some  of  the  solid  unknown  with  dilute  sulphuric  acid. 
An  evolution  of  gas  indicates  : 

HCN  —  Noted  by  odor  of  peach  blossoms  (caution  !).     Belongs 

in  Group  II. 

NO2   —  Indicates  nitrites ;  noted  by  color,  odor,  solubility. 
H2S    •  -  Noted  by  odor  and  by  lead  acetate  paper. 
SO2    —  With    separation    of     sulphur  —  from    thiosulphates. 

Noted  by  odor  and  reduction  of  permanganic  acid. 
SO2    — Without  separation  of  sulphur  —  from  sulphites. 
CO2    —  From  carbonates ;  noted  by  turbidity  produced  when 

gas  is  passed  into  limewater. 
C12      —  From  hypochlorites  ;  noted  by  odor  and  other  tests. 

On  boiling  with  dilute  acid,  if  no  reaction  has  occurred  in  the 
cold,  or  when  action  has  ceased,  there  may  be  formed : 

CO2  —  From  certain  carbonates  ;  noted  as  above. 

O2  —  From  peroxides  ;  noted  by  support  of  combustion. 

May  be  formed  in  the  cold  also. 

HCN  —From  ferro-  or  ferricyanides.     See  Group  II. 

HC2H3O2  —  From  acetates;  noted  by  odor  of  vinegar. 

If  the  residue  is  colored,  it  may  contain  certain  sulphides,  and 
these  may  be  sought  by  reaction  I  or  4,  page  104. 

If  any  of  the  indications  are  observed,  separate  tests  are 


ACID  ANALYSIS  1 07 

made  for  all  suspected  acids,  according  to  the  preliminary  tests. 
See  also  exercises. 

A  portion  of  the  unknown  is  dissolved  in  water l  and  treated 
with  hydrochloric  acid,  until  completely  decomposed,  evaporated 
to  dryness  and  heated  gently,  and  redissolved  as  completely  as 
possible  in  water  with  a  few  drops  of  HC1.  An  insoluble 
residue  is  SiO2  and  may  be  confirmed  by  3,  p.  115. 

EXERCISES 

I.  If  sulphides  and  carbonates  are  present  together,  how  are 
they  to  be  distinguished  ? 

II.  If  on  testing  with  lead  acetate  paper  for  sulphides  a  yel- 
low color  is  produced  upon  the  paper,  what  is  present  ? 

III.  If  the  gases  from  a  sample  treated  with  sulphuric  acid 
are  passed  into  acidified  bichromate  solution  with  the  formation 
of  a  green  color,  what  gases  may  be  present  ? 

IV.  If  sulphites  and  sulphides  were  both  present  in  the  same 
sample,  what  substance  would  be  formed  on  acidification  ? 

V.  Make  a  list  of  the  important  sulphide  minerals. 

VI.  What    cyanides    are    commercially    important?      Give 
methods  of  preparation  of  four  different  cyanides. 

Group  II 
HC1,  HBr,  HI,  HCN,  H3Fe(CN)6,  H4Fe(CN)6. 

Preliminary  Reactions 
HYDROCHLORIC  ACID 

1.  Dilute  sulphuric  acid  produces  no  evidence  of  reaction. 

2.  Concentrated  sulphuric  acid  on  the  dry  salt  gives  colorless 
HC1,  fuming  in  air. 

1  The  unknowns  consist  wholly  of  salts  which  are  soluble  in  water.  If  samples  which 
are  not  soluble  are  encountered,  they  are  prepared  for  examination  as  described  on 
p.  no.  The  nitrate  will  contain  sodium  silicate  if  insoluble  metasilicates  were  originally 
present. 


IO8  QUALITATIVE  ANALYSIS 

3.  Silver  nitrate  produces  a  curdy  white  precipitate  of  silver 
chloride;  soluble  in  ammonium  hydroxide,  potassium  cyanide, 
sodium  thiosulphate,  and  ammonium    '  sesquicarbonate.'     (See 
Appendix.) 

4.  The  dry  salt  treated  with  concentrated  sulphuric  acid  and 
potassium  dichromate  gives  a  brownish  vapor  of  chromium  oxy- 
chloride.     If  this  vapor  is  dissolved  in  water,  it  produces  chromic 
acid.     The  chromic  acid  may  be  detected  by  the  hydrogen  per- 
oxide reaction.     See  chromic  acid.     (Distinction  from  HI  and 

HBr.) 

HYDROBROMIC  ACID 

1.  Concentrated  sulphuric  acid  on  the  dry  salt  gives  rise  to 
colored. gases  (distinction  from  HC1)  which  fume  in  the  air  arid 
do  not  render  water  held  on  a  glass  rod  turbid  (distinction  from 
HF). 

2.  Silver  nitrate  produces  a  curdy  yellowish  precipitate ;  sol- 
uble in  ammonium  hydroxide,  potassium  cyanide,  sodium  thio- 
sulphate, but  not  in  ammonium  'sesquicarbonate.' 

3.  Chlorine  water  liberates  bromine  ;  soluble  in  chloroform  or 
carbon  disulphide,  changed  by  excess  to  yellowish  BrCl  (distinc- 
tion from  HI). 

4.  Potassium  dichromate  liberates  no  bromine  in  dilute  sul- 
phuric acid  solution  (difference  from  HI). 

5.  Potassium  dichromate  and  concentrated  sulphuric  acid  on 
the  dry  salt  liberate  bromine. 

6.  Nitrites  in  acid  solution,  if  dilute,  free  no  bromine  in  the 
cold. 

HYDRIODIC  ACID 

1.  Concentrated  sulphuric  acid  reacts  in  the  cold,  liberating 
hydriodic  acid  and  iodine  and  sulphur  dioxide,  hydrogen  sul- 
phide, and  sulphur. 

2.  Silver  nitrate  produces  a  curdy  yellow  precipitate ;  soluble 
in  potassium  cyanide  and  sodium  thiosulphate  ;  but  only  slightly 


ACID  ANALYSIS  109 

soluble  in  ammonium  hydroxide  and  ammonium  '  sesquicarbon- 
ate '  (difference  from  HC1  and  HBr). 

3.  Lead  salts  precipitate  yellow  lead  iodide. 

4.  Cupric  salts  produce  a  separation  of  cuprous  iodide  and 
iodine,  made  nearly  white  by  sulphurous  acid. 

5.  Nitrites  in  dilute  acid  solution  liberate  iodine. 

6.  Chlorine  water  liberates  iodine  which  colors  starch  emul- 
sion blue  or  chloroform  violet,  in  excess  forms  colorless  solutions 
of  iodates. 

7.  Potassium    dichromate  and  both  concentrated  and  dilute 
sulphuric  acid  liberate  iodine;  detected  by  chloroform  or  carbon 
disulphide. 

HYDROCYANIC  ACID  (PRUSSIC  ACID) 

1.  Silver  nitrate  precipitates  silver  cyanide;   soluble  in  excess 
of  the  alkali  cyanide. 

2.  Ammonium    sulphide   boiled   with   alkali   cyanides  gives 
sulphocyanates.      This    solution,    if    made    slightly    acid,    and 
heated  until  a  drop  gives  no  black  precipitate  with  lead  salts, 
will  give  a  blood-red  color  with  ferric  chloride. 

3.  Addition  of  a  small  amount  of  ferrous  sulphate  and  an 
excess  of  sodium  hydroxide  will,  on  heating,  form  potassium 
ferrocyanide.     This  solution,  made  acid  with  hydrochloric  acid 
and  treated  with  ferric  chloride,  leaves  an  insoluble  precipitate 
of  Prussian  blue. 

4.  Sulphuric  acid  liberates  prussic  acid  in  the  cold,  which 
may  be  noted  by  the  odor  of  peach  blossoms  (virulent  poison ! 
caution !). 

HYDROFERROCYANIC  ACID 

1.  Silver  nitrate  produces  a  white  precipitate  of  silver  ferro- 
cyanide. 

2.  Ferric  salts  produce  a  precipitate  of  Prussian  blue.     (See 
Treadwell  and  Hall,  p.  100.) 


HO  QUALITATIVE  ANALYSIS 

3.  With  exclusion  of  air,  ferrous  salts  produce  a  white  precipi- 
tate, which  rapidly  becomes  blue  by  oxidation. 

HYDROFERRICYANIC  ACID 

1.  Silver  nitrate  produces   an    orange    precipitate  of   silver 
ferricyanide. 

2.  Ferric  salts  produce  a  brown  coloration. 

3.  Ferrous  salts  produce  a  precipitate  of  Turnbull's  blue. 
The  analysis  of  samples  3  and  4  of  set  of  unknowns  is  now 

undertaken,  using  the  following  scheme  as  a  guide. 

Analysis  of  Group  II 

For  preparation  of  a  solution  of  the  anions,  see  p.  115. 

Acidify  the  solution  with  nitric  acid,  adding  a  slight  excess, 
and  treat  with  silver  nitrate.  If  no  precipitate  is  formed,  the 
members  of  this  group  are  not  present.  If  a  precipitate  forms, 
a  fresh  portion  of  the  solution  is  acidified  with  hydrochloric  acid 
and  a  few  drops  of  ferrous  sulphate  are  added;  an  immediate 
blue  precipitate  indicates  ferricyanic  acid ;  a  white  one,  rapidly 
turning  blue,  indicates  ferrocyanic  acid.  Each  of  these  is  then 
sought  by  other  of  the  preliminary  reactions.  If  either  is 
present,  in  order  to  detect  hydrocyanic  acid  some  of  the  dry 
sample  is  placed  in  a  porcelain  dish,  treated  with  dilute  hydro- 
chloric acid,  and  covered  with  a  watch  glass,  which  is  moistened 
with  ammonium  sulphide,  and  allowed  to  stand  for  some  time. 
The  cover  is  rinsed  off  with  water  and  the  rinsing  acidified 
and  treated  with  a  drop  or  two  of  ferric  chloride ;  a  blood-red 
color  indicates  hydrocyanic  acid  in  the  original  substance.  If 
any  of  the  cyanogen  compounds  are  present,1  a  larger  portion 
of  the  solution  is  treated  with  an  excess  of  nickel  sulphate  and 

1  If  ferricyanides  or  ferrocyanides  are  present,  it  must  be  observed,  that  on  acidifying 
the  solution  with  nitric  acid,  hydriodic  acid  is  destroyed.  In  such  case  the  presence  of 
iodine  or  its  salts  may  be  detected  on  the  dry  substance  by  reaction  I,  page  108. 


ACID  ANALYSIS  III 

filtered ;  the  filtrate  contains  only  the  nickel  salts  of  hydro- 
chloric, hydrobromic,  and  hydriodic  acids,  and  sulphates.  The 
nickel  is  removed  by  precipitation  with  sodium  hydroxide  and 
the  filtrate  is  divided  into  two  portions.  To  one  portion  is 
added  a  drop  of  chloroform  and  a  drop  of  chlorine  water  and 
shaken  ;  a  violet  color  indicates  iodine.  If  present,  remove  by 
boiling  with  potassium  dichromate  and  dilute  sulphuric  acid  as 
long  as  fumes  of  iodine  are  developed,  and  test  filtrate  with 
chlorine  water  for  bromide.  (If  very  small  amounts  of  iodides 
are  present,  they  may  be  removed  by  an  excess  of  chlorine ;  and 
if  bromine  is  present,  the  brown  color  will  be  imparted  to  the 
chloroform.  If  very  large  quantities  of  iodides  are  present, 
they  may  be  removed  by  precipitation  with  copper  sulphate  and 
the  test  for  bromides  then  made  in  the  filtrate.)  To  the  other 
portion  silver  nitrate  is  added  and  the  precipitate  is  treated, 
after  washing,  with  ammonium  sesquicarbonate  and  boiled  and 
filtered;  the  filtrate  is  acidified  with  nitric  acid,  and  a  white 
precipitate  indicates  chlorides.  (The  presence  of  chlorides  may 
also  be  shown  by  reaction  4,  p.  108.) 

The  residue  insoluble  in  ammonium  sesquicarbonate  is 
treated,  on  the  filter,  with  dilute  ammonium  hydroxide  and 
the  filtrate  acidified;  a  yellowish  precipitate  indicates  bromides; 
an  insoluble  residue  on  the  filter  indicates  iodides. 

A  more  delicate  and  in  some  respects  more  satisfactory 
method  of  separation  of  the  halogens  is  to  acidify  the  filtrate 
after  the  removal  of  the  cyanides  and  nickel,  and  boil  with 
ferric  sulphate  or  ferric  alum,  an  evolution  of  violet  vapors 
indicating  iodine.  When  this  reaction  is  over,  the  solution 
is  boiled  with  a  crystal  of  potassium  permanganate,  and  when 
the  removal  of  the  bromine  is  thus  effected,  the  solution  is 
filtered  and  diluted  (to  prevent  precipitation  of  silver  sulphate) 
and  treated  with  silver  nitrate,  a  white  precipitate  of  silver 
chloride  indicating  the  presence  of  chlorides. 


112  QUALITATIVE  ANALYSIS 

EXERCISES 

I.  Make   a   list   of   the   mineral   chlorides   which   are   com- 
mercially important,  either  for  their  own  sake  or  for  the  sake 
of  the  substances  prepared  from  them. 

II.  In  the  analysis  of  the  group  by  the  method  given,  why 
is  nickel  removed,  and  why  not  use  ammonium  hydroxide  for 
the  purpose?     See  reaction  No.  215. 

III.  If  the  halides  were  present  in  the  original  substance  as 
silver  salts,  devise  or  look  up  a  method  of  separation. 

IV.  Look  up  and  make  a  synopsis  of   Remsen's  theory  of 
double  salts  of  the  halides.      See  Remsen's  Advanced  Course, 
p.  465. 

Group  III 
Preliminary  Reactions 

H2SO4,    H2CrO4,    H3AsO4,    H3AsO3,    H3PO4,    H3BO3,    HF, 
H4Si04,  H2Si03,  H2C204,  H2C4H406 

SULPHURIC  ACID 

1.  Silver  nitrate  precipitates  silver  sulphate  from  sufficiently 
concentrated  solutions.     (Solubility  of  silver  sulphate,  0.58  g. 
per  100  c.c.) 

2.  Barium  chloride  precipitates  barium  sulphate  ;  practically 
insoluble  in  all  acids. 

3.  Lead  acetate  precipitates  lead  sulphate ;    soluble  in  con- 
centrated sulphuric  acid,   ammonium   acetate,   ammonium    tar- 
trate. 

CHROMIC  ACID 

1.  All    chromates    and    dichromates    are    colored,    and    are 
reduced   in    acidified  solution    by   hydrogen  sulphide,   sulphur 
dioxide,  alcohol,  and  hydriodic  acid,  forming  green  solutions. 

2.  Silver  nitrate  precipitates  silver  chromate  from  neutral  or 
slightly  acid  solutions. 


ACID  ANALYSIS  113 

3.  Lead    acetate    precipitates   lead    chromate,  in    solutions 
acidified  with  acetic  acid,  in  presence  of  ammonium  acetate. 

4.  Hydrogen   peroxide   added    to    dilute   cold    solutions   of 
chromic  acid,  or  acidified  chromates,  produces  a  blue  color,  due 
to  the  formation  of  perchromic  acid.     This  substance  is  ren- 
dered more  stable  by  solution  in  ether.     The  more  dilute,  and 
the   lower   the  temperature,   the   more   easily  is   the   reaction 
carried  out. 

ARSENIC  ACID 

1 .  Silver  nitrate  precipitates  from  neutral  solutions  chocolate- 
brown  silver  arsenate;  soluble  in  acid  and  in  ammonia  (differ- 
ence in  color  from  arsenous  and  phosphoric  acids). 

2.  Hydrogen    sulphide    precipitates    arsenic    solutions    only 
slowly.     In  hot,  strongly  acid  solutions  it  precipitates  a  mixture 
of  the  tri-  and  pentasulphides. 

3.  Magnesia    mixture    precipitates    magnesium    ammonium 
arsenate,  white  and  crystalline. 

4.  Ammonium   molybdate   in   nitric   acid   solution   (see   ap- 
pendix) precipitates  yellow  arsenomolybdate. 

ARSENOUS  ACID 

1.  Silver  nitrate  in  neutral  solutions  precipitates  yellow  silver 
arsenite  (difference  in  color  from  arsenic  acid). 

2.  Hydrogen  sulphide  precipitates  arsenous  arsenic  from  acid 
solutions  at  once  as  the  trisulphide.     In  the  absence  of  other 
salts  the  arsenous  sulphide  may  be  colloidal.     For  colloidal  so- 
lutions  and   their   treatment  see  Ostwald's   Scientific  Founda- 
tions, p.  25. 

3.  Iodine  solutions  are  decolorized  by  solutions  alkaline  with 
sodium  carbonate,  arsenates  being  formed. 

4.  Magnesia  mixture  and  ammonium  molybdate  produce  no 
precipitate. 


114  QUALITATIVE  ANALYSIS 

PHOSPHORIC  ACID 

1.  Silver  nitrate  produces  in  neutral  solution  a  yellow  pre- 
cipitate of  the  orthophosphate ;   soluble  in  nitric  acid  and  in 
ammonium  hydroxide.      (The   meta-   and    pyrophosphates   are 
white.     For  other  distinctions,  see  Treadwell  and  Hall.) 

2.  Hydrogen  sulphide  produces  no  visible  effect  on  solutions 
of  the  acid  or  of  its  salts. 

3.  Magnesia  mixture  produces  a  precipitate  of  magnesium 
ammonium  phosphate. 

4.  Ammonium  molybdate  on  standing  produces  a  quantita- 
tive precipitate  of  ammonium  phosphomolybdate ;  readily  solu- 
ble in  the  alkalies. 

5.  Ferric  chloride  and  sodium  acetate  produce  a  quantitative 
precipitate  of  ferric  phosphate.     (A  large  excess  of  ferric  chlo- 
ride is  to  be  avoided.)     (Why  ?) 

BORIC  ACID  (SOLUTION  OF  BORAX  FOR  REACTIONS) 

1.  Silver  nitrate  in  neutral  solutions  precipitates  silver  meta- 
borate,  changed  on  heating  to  silver  oxide.      Dilute  solutions 
give  the  oxide  directly.     (Why  ?) 

2.  Mercuric  chloride  produces  a  red  precipitate. 

3.  Concentrated  sulphuric  acid  and  alcohol  (methyl  alcohol 
preferred),  heated  with  borax  and  ignited,  produce  a  green  flame, 
due  to  the  combustion  of  the  volatile  methyl  orthoborate. 

4.  Turmeric  paper,  dipped  in  a  solution  of  free  boric  acid  and 
dried  at  not  over  100°,  is  turned  rose-red,  and  the  color  is  not 
removed  by  treatment  with  acids.     Ammonia  turns  it  green  or 
blue.     Borax  acidified  with  hydrochloric  acid  produces  the  same 
result.     (For  discussion,  see  Treadwell  and  Hall,  p.  311.) 

HYDROFLUORIC  ACID  (SODIUM  FLUORIDE  FOR  REACTIONS) 

i.    Barium  chloride  precipitates  a  flocculent,  bulky  precipitate 
from  solutions  of  the  soluble  fluorides. 


ACID   ANALYSIS  115 

2.  All  fluorides  when  heated  with  concentrated  sulphuric  acid 
give  rise  to  hydrofluoric  acid,  which  acts  on  glass,  or  other  sili- 
cates, if  present,  and  forms  silicon  tetrafluoride,  which  in  turn  is 
decomposed  by  water  with  the  formation  of  fluosilicic  and  ortho- 
silicic  acids.  This  reaction  is  most  easily  carried  out  by  placing 
the  fluoride  and  acid  in  a  test  tube,  and,  while  holding  near  the 
surface  of  the  mixture  a  glass  rod  with  a  drop  of  water,  warming 
gently.  In  the  presence  of  a  fluoride,  a  crust  of  silicic  acid  is 
formed  on  the  water  drop. 

SILICIC  ACID  (SEE  AI/SO  GROUP  I) 

1.  See  p.  104. 

2.  Silicates  insoluble  in  water  are  converted  into  soluble  so- 
dium silicate  by  fusion  on  platinum  with  sodium  carbonate. 

3.  Silicates  treated  on  platinum  or  lead  with  hydrofluoric  acid 
are  converted  into  silicon  tetrafluoride,  which  reacts  with  a  drop 
of  water  (using  a  platinum  wire,  with  a  loop  to  hold  the  water 
drop).     See  reaction  2  on  HF. 

4.  Mineral  silicates  are  usually,  but  not  always,  converted  to 
soluble  sodium  silicate  by  boiling  with  sodium  carbonate. 

5.  Barium  salts  precipitate  crystalline  barium  silicate. 

OXALIC  ACID  (SEE  GROUP  V) 
TARTARIC  ACID  (SEE  GROUP  V) 

The  analysis  of  samples  5  and  6  of  set  of  unknowns  is  now 
undertaken,  using  the  following  scheme  as  a  guide. 

Analysis  of  Group  III 

A  solution  is  prepared  by  boiling  the  sample  with  a  saturated 
solution  of  sodium  carbonate  (i  gram  of  the  sample  to  15  c.c.  of 
the  solution)  for  fifteen  minutes,  replacing  the  water  evaporated. 
The  hot  solution  is  diluted  with  an  equal  quantity  of  water  and 


Il6  QUALITATIVE  ANALYSIS 

filtered.1  This  solution  is  made  slightly  acid  with  hydrochloric 
acid  and  neutralized,  drop  by  drop,  with  ammonium  hydroxide, 
using  litmus  as  an  indicator.  A  portion  is  treated  with  a  few 
drops  of  calcium  and  barium  chloride  solutions.  (The  calcium 
salt  is  added  because  the  calcium  salts  of  some  of  the  group  are 
more  insoluble  than  the  barium  salts.)  If  no  precipitate  is 
formed,  the  members  of  this  group  are  absent  (except  boric  acid, 
the  barium  and  calcium  salts  of  which  are  fairly  soluble). 

Test  is  made  on  the  original  solid  material  for  boric  acid  by 
reaction  3,  page  114,  and  for  silicic  acid  by  reaction  3,  page  115. 

If  a  precipitate  is  formed,  then  tests  for  the  other  acids  are 
made  separately,  using  by  preference  reactions  2  and  3  for  sul- 
phuric acid,  reactions  I  and  4  for  chromic  acid  (if  the  solution  is 
colorless,  chromic  acid  is  absent),  reactions  2  and  4  for  arsenic 
acid,  reactions  3  and  4  for  arsenous  acid.  If  either  of  these 
latter  is  present,  it  must  be  removed  by  hydrogen  sulphide  and 
the  solution  boiled  to  expel  hydrogen  sulphide  before  testing 
for  phosphoric  acid,  using  reactions  2,  3,  and  4,  p.  1 14. 

For  oxalic  and  tartaric  acids,  see  analysis  of  Group  V. 

EXERCISES 

I.  An  insoluble  residue  nearly  always  remains  after  treatment 
with  sodium  carbonate.     What  substances  may  it  contain  ? 

II.  In  case  the  substance  to  be  analyzed  is  soluble  in  water 
and  the  treatment  with  sodium  carbonate  is  omitted,  what  sub- 
stances may  be  in  the  precipitate  which  are  not  representative  of 
this  group  ? 

III.  Give  an  outline  of  the  manufacture  of  "  superphosphate." 

IV.  What  forms  of  boron  are  found  in  nature  and  where  ? 

1  The  residue  usually  consists  wholly  of  carbonates  and  should  dissolve  completely  in 
dilute  HC1.  Any  portion  which  remains  undissolved  in  the  acid  may  be  sulphates  or 
silicates  of  difficultly  decomposed  minerals  and  may  be  fured  with  sodium  carbonate  in  iron, 
or  preferably  platinum,  and  the  aqueous  solution  of  the  fused  mass  tested  for  sulphates  and 
silicates. 


ACID  ANALYSIS  Ii; 

V.  In  case  of  an  alloy  containing  phosphorus  or  sulphur,  what 
treatment  is  needed  to  detect  them  ? 

VI.  Why  is  it  necessary  to  acidify  arsenic  solutions  in  order 
to  precipitate  arsenic  sulphide  with  hydrogen  sulphide  ? 

Group  IV 
HN03,  HC103,  HMn04,  HC2H3O2 

Preliminary  Reactions 

NITRIC  ACID 

1.  Ferrous  sulphate  is  oxidized  by  concentrated  nitric  acid  to 
ferric  salts,  and  nitric  oxide  is  formed,  which  unites  with   an 
excess  of  ferrous  sulphate  to  form  a  brown,  unstable  compound, 
riitrosyl  ferrous  sulphate,  which  decomposes  on  boiling.     The 
simplest  method  of  using  the  test  is  to  add  to  the  solution  to  be 
tested  a  few  cubic  centimeters  of  ferrous  sulphate  solution  and 
pour  down  the  side  of  the  containing  test  tube  3-5  cubic  centi- 
meters of  concentrated  sulphuric  acid  ;  a  brown  ring  will  appear 
at  the  junction  of  the  two  liquids.     If  organic  acids  are  present 
and  are  charred  by  the  concentrated  sulphuric  acid,  the  color 
will  not  disappear  on  boiling.     If  the  solution  is  acid,  it  is  best 
to  add  a  few  cubic  centimeters  of  the  concentrated  acid,  cool, 
and  then   superimpose  a   layer   of   ferrous  sulphate   solution. 
Iodides,    bromides,    ferrocyanides,    chromates,    permanganates, 
and  chlorates  interfere  with  this  test. 

2.  A  few  drops  of  solution  of  nitrates  evaporated  on  a  porce- 
lain dish  just  to  dryness  and  gently  warmed  with  a  few  drops 
of  phenolsulphuric  acid  form  picric  acid,  which,  when  cooled 
and  made  alkaline  with  ammonium  hydroxide,  gives  an  intense 
yellow  color. 

3.  One  cubic  centimeter  of  brucine  solution  added  to  a  mix- 
ture of  i  part  of  the  sofution  and  3  parts  concentrated  sulphuric 
acid  produces  a  deep  red  color,  changing  to  orange,  then  to  yel- 


Il8  QUALITATIVE  ANALYSIS 

low,  and  finally  to  greenish  yellow.     (Nitrous  acid  does  not  give 
this  test.) 

4.  Iodine  is  not  liberated  from  potassium  iodide  solutions  by 
dilute  nitric  acid  (see  Nitrous  Acid). 

CHLORIC  ACID 

1.  Concentrated  sulphuric  acid    decomposes  chlorates  with 
liberation  of  the  greenish  yellow  chlorine  dioxide  (caution,  ex- 
plosive) and  chlorine. 

2.  Concentrated  hydrochloric  acid  decomposes  all  chlorates 
with  the  liberation  of  chlorine  (euchlorine,  see  Remsen,  Ad- 
vanced Course,  p.  1 2 1 ). 

3.  Reducing  agents  decompose  chlorates  in  neutral,  acid  or 
alkaline  solutions,  e.g.  nascent  hydrogen. 

4.  Ferrous  salts  are  oxidized  by  chlorates  to  ferric  salts  on 
boiling  with  dilute  sulphuric  acid. 

PERMANGANIC  ACID 

1.  All  permanganates  are  soluble  in  dilute  acids  and  are  re- 
duced to  manganese  dioxide  in  alkaline  solution  and  to  manga- 
nous  salts  in  acid  solution  by  reducing  agents,  SnCl2,  H2SO3,  etc. 

2.  Hydrogen  peroxide  in  acid  solution  decolorizes  permanga- 
nates with  a  lively  evolution  of  oxygen. 

ACETIC  ACID  (See  GROUP  V) 

Analyses  of  samples  7  and  8  are  now  undertaken  according 
to  the  following  scheme. 

Analysis  of  Group  IV 

If  the  aqueous  solution  is  colored  red  by  permanganic  acid, 
it  is  decolorized  by  hydrogen  peroxide  and  a  portion  evaporated 
to  dryness  on  the  water  bath  and  tested  for  chlorates  by  re- 
action i.  If  chlorates  are  present,  a  second  portion  is  treated 


ACID  ANALYSIS  119 

with  sulphur  dioxide  while  being  boiled.     After  cooling,  it  is 
tested  for  nitric  acid  by  reactions  I,  2,  or  3  for  nitric  acid. 

A  third  portion  is  acidified  and  treated  with  hydrogen  sulphide 
until  the  oxidizing  acids  are  removed;  filtered  and  evaporated 
nearly  to  dryness  and  treated  for  acetic  acid  as  in  Group  V. 

EXERCISES 

I.  Devise  a  method  for  removing  the  acids  which  interfere 
with  the  ferrous  sulphate  test  for  nitric  acid. 

II.  If  dry  permanganates  are  treated  with  concentrated  sul- 
phuric acid,  what  is  formed,  and  why  is  the  operation  dangerous  ? 

III.  How  is  potassium  permanganate  made? 

IV.  How  are  chlorates  made  ? 

V.  Look  up  the  history  of  euchlorine  and  state  what  is  its 
bearing  on  the  demonstration  of  the  elementary  nature  of  chlo- 
rine.    See  Alembic  Club  reprints,  No.  13. 

Group  V 

H2C2O4,  H2C4H4O6,  HC2H3O2,  etc. 

Preliminary  Reactions 

OXALIC  ACID 

1.  Concentrated  sulphuric  acid  heated  with  an  oxalate  evolves 
carbon  monoxide  and  carbon  dioxide.     The  former  is  detected 
by  passing  the  mixed  gases  through  baryta  water  or  sodium  hy- 
droxide and  igniting  the  insoluble  monoxide. 

2.  Calcium  salts,  even  calcium  sulphate,  precipitate  oxalates ; 
soluble  in  hydrochloric,  but  not  in  acetic  acid. 

3.  Silver    nitrate    precipitates    white    curdy    silver   oxalate; 
soluble  in  ammonium  hydroxide  and  dilute  acids. 

4.  Heat  decomposes  the  oxalates  with  but  slight  carbonization, 
some  of  them  leaving  carbonates  of  the  metals  and  others  the 
metals  themselves. 


120  QUALITATIVE  ANALYSIS 

TARTARIC  ACID 

1.  Ignition  of  the  tartrates  leaves  a  residue  of  carbon,  gives 
an  odor  of  burnt  sugar,  and  leaves  either  a  carbonate  of  the 
metals  or  the  metals  themselves.     (See  Treadwell  and  Hall, 

P-  3170 

2.  Concentrated  sulphuric  acid,  when  heated  with  tartrates, 
gives  carbon   monoxide,  carbon  dioxide,  and  sulphur  dioxide. 
The  gases  may  be  passed  through   acidified  bichromate,  then 
through  alkaline  solution,  and  the  issuing  gas  burned. 

3.  Barium  chloride  precipitates  barium  tartrate. 

ACETIC  ACID 

1.  Concentrated  sulphuric  acid  sets  acetic  acid  free  from  its 
salts,  and,  if  alcohol  is  added  at  the  same  time,  ethyl  acetate  is 
produced,  which  may  be  noted  by  its  "  fruity  "  odor. 

2.  Ignition  decomposes  all  acetates,  leaving  behind  oxides, 
carbonates,  or  metals,  with  the  evolution  of  combustible  gases, 
but  not  always  with  the  formation  of  free  carbon. 

3.  Ferric  chloride  in  neutral  solutions  produces  a  blood-red 
ferric  acetate,  decomposed  by  boiling  into  a  reddish  precipitate 
of  basic  ferric  acetate. 

OTHER  ORGANIC  ACIDS 

Nearly  all  other  organic  acids  char  on  ignition  and  leave 
carbonates  of  the  alkali  metals  when  such  are  present.  These 
acids  are  not  usually  present  in  inorganic  mixtures,  and  their 
discussion  will  not  form  a  part  of  this  treatment. 

No  unknown  samples  are  given  for  Group  V,  since  the  acids 
which  are  here  presented  are  found  also  in  Groups  III  and  IV. 
When  organic  matter  is  present,  the  search  for  these  may  be 
conducted  as  given  below. 


ACID  ANALYSIS  121 

Analysis  of  Group  V 

A  portion  of  the  solution  to  be  tested  for  members  of  this 
group  may  be  treated  with  a  solution  of  calcium  sulphate,  and 
a  precipitate  soluble  in  hydrochloric  acid  without  evolution  of 
carbon  dioxide  indicates  the  presence  of  oxalic  acid.  The  re- 
mainder is  evaporated  to  near  dryness  on  the  water  bath,  and  a 
portion  may  be  tested  for  tartaric  acid  by  reaction  2. 

Another  portion  may  be  treated  for  acetic  acid  as  in  reaction 
I,  or  it  may  be  made  exactly  neutral  and  treated  with  ferric 
chloride  as  in  reaction  3. 

EXERCISES 

I.  How  is  sugar  of  lead  manufactured  ?     For  what  purpose 
is  it  used  ? 

II.  What  is  "  cream  of  tartar,"  and  how  is  it  purified  ? 

III.  How  is  oxalic  acid  made  commercially? 

IV.  What  is  pyroligneous  acid,  and  what  substances  may  be 
obtained  from  it  ? 

When  the  student  is  made  somewhat  familiar  by  a  study  of 
the  preceding  reactions  and  the  twenty  (or  more,  if  desired) 
samples  containing  members  of  single  groups  are  analyzed,  he 
may  be  considered  ready  to  undertake  the  complete  analysis  of 
unknown  substances.  This  involves  all  the  steps  already  given, 
and  some  modifications  which  are  given  in  Part  IV.  The 
directions  for  group  separations  are  given  in  abbreviated  and 
tabular  form  for  convenience  in  use. 


PART    IV 

SYSTEMATIC  ANALYSIS 

The  complete  analysis  of  a  substance  may  be  considered  as 
involving  three  steps : 

1.  Preliminary  examination. 

2.  Examination  for  metals  (cations). 

3.  Examination  for  non-metals  (anions). 

A  preliminary  examination  is  needed  in  order  to  proceed  in- 
telligently to  get  the  substance  into  solution.  It  also  serves  to 
give  information  which  renders  subsequent  examination  more 
accurate  and  interesting,  and  at  times  renders  such  analysis 
brief  or  even  unnecessary.  It  is  also  of  value  where  the  pres- 
ence of  some  one  element  is  to  be  determined  and  recourse  to  a 
complete  analysis  is  not  desirable.  It  is  not  essential  that  all 
the  steps  in  the  following  outline  be  rigidly  followed  in  every 
case  when  a  complete  examination  is  contemplated,  but  they 
are  to  be  followed  so  far  as  to  enable  the  worker  to  proceed  in- 
telligently in  the  later  work.  The  closed  tube  test  is,  however, 
never  to  be  omitted,  since  it  serves  to  detectahe  presence  of 
organic  matter  which  must  always  be  removed  before  a  mineral 
analysis  is  undertaken.  The  sample  is  first  powdered,  and  then 
small  portions  of  the  homogeneous  "sample  are  taken  for  the 
separate  tests.1  As  a  general  rule,  it  is  wise  to  reserve  at  least 
one  half  of  the  sample.  Of  the  other  half,  one  third  is  reserved 
for  the  metal  analysis,  one  third  for  the  acid  analysis,  and  the 
remaining  portion  used  for  the  preliminary  examination. 

1  If  the  substance  is  a  liquid  or  alloy,  see  p.  127. 

122 


SYSTEMATIC  ANALYSIS 


123 


Closed  Tube  Test 

A  hard  glass  test  tube  is  to  be  used  or  a  small  tube  sealed 
at  one  end.  A  small  amount  of  the  substance  is  placed  in  the 
clean  and  dry  tube,  and  heated  gently  at  first  and  then  to  the 
highest  heat  of  the  Bunsen  flame.  Careful  observation  is  made 
of  any  visual  phenomena,  and  the  odor  and  combustibility  of 
the  issuing  gases  are  tested  from  time  to  time ;  the  latter  is  best 
done  by  means  of  a  match  or  glowing  splinter.  The  phenomena 
which  may  occur  and  their  interpretation  are  given  below : 


PHENOMENA 


INDICATION 


1 .  Moisture  without  carbonization. 

2.  Carbonization  with  odor  and  com- 

bustible gases. 

3.  Sublimate. 

(a)  Black,  with  garlic  odor. 

(b)  Black     or    gray,     forming 

globules  when  rubbed,  or 
forming  a  mirror. 

(c)  Black,    turning    red    when 

rubbed. 

(d)  Black,  with  violet  vapors. 

(e)  Reddish  brown  vapors,  cool- 

ing to  a  yellow  solid. 
(/)  Yellow  to  red. 
(g)  White,  easily  volatile. 

(h)  White,  volatile   with   diffi- 
culty. 

4.  Gases. 

(a)  Kindles  a  glowing  splinter. 

(£)  Brownish  red,  with  charac- 
teristic odor. 

(c)  Colorless,    with    character- 
istic odor. 
d    Characteristic  odor. 


Water. 
Organic  matter. 

Arsenic. 

Mercury. 

Mercury  sulphide. 

Iodine.  , 

Sulphur  or  persulphides. 
Sulphur  or  arsenic  sulphide. 
Arsenious   oxide,   calomel,   corrosive 

sublimate,  or  ammonium  salts. 
Antimony  oxide. 


Oxygen  —  from  peroxides,  chlorates, 
or  nitrates. 

Nitrogen  peroxide  —  from  nitrates  or 
nitrites. 

Sulphur  dioxide  —  from  sulphates,  sul- 
phides, or  sulphites. 

Ammonia  —  from  ammonium  salts  or 
a  cyanide. 


124 


QUALITATIVE  ANALYSIS 


PHENOMENA 


INDICATION 


(e)  Characteristic  odor,  black- 
ening lead  acetate  paper. 
(/)  Burning  with  a  blue  flame. 


(g)  Causing  turbidity  in   drop 

of  limewater. 
^    Colored  —  characteristic. 


Hydrogen  sulphide  —  from  moist  sul- 
phides. 

Carbon  monoxide  —  from  an  oxalate 
or  tartrate,  or  methane  from  an 
acetate. 

Carbon  dioxide  —  from  carbonates  or 
organic  matter. 

Chlorine,  bromine,  iodine. 


5.    Non-volatile  residue  may  change  color ;  if  so,  the  indications  are  as  follows  : 


ORIGINAL  COLOR 

COLOR  HOT 

COLOR  COLD 

INDICATION 

White 

Yellow 

White 

Zinc  oxide 

White  or  yellow 
White  or  yellow 
White  or  yellow 
Yellow  or  brown 

Yellow 
Brownish  red 
Brownish  red 
Black 

Pale  yellow 
Deep  yellow 
Pale  yellow 
Brownish  red 

Tin  oxide 
Lead  oxide 
Bismuth  oxide 
Iron  oxide 

Yellow  or  red 

Green 

Green 

Chromium 

Pink,  green,  or  blue 

Black 

Black 

Nickel,  copper,  cobalt 

6.    The  substance  melts  and  remains  liquid  with  or  without  expulsion  of  vapors 
—  alkali  salts. 

The  Bead  Test 

A  bead  of  borax  is  made  on  a  loop  at  the  end  of  a  straight 
platinum  wire,  and  after  cooling,  is  moistened  with  the  tip  of  the 
tongue,  and  a  very  little  of  the  solid  substance  to  be  tested  is 
caused  to  adhere  to  the  bead,  and  the  whole  introduced  into  the 
oxidizing  flame  of  the  Bunsen  burner.  After  the  color  of  the 
transparent  bead  is  observed,  it  is  brought  into  the  reducing 
flame,  and  after  some  minutes  the  color  is  again  observed.  The 
results  and  their  indications  follow  : 


OXIDIZING  FLAME 

REDUCING  FLAME 

INDICATION 

Blue 

Blue 

Cobalt 

Green 

Green 

Chromium 

Greenish  blue 
Amethyst 
Brownish  red 
Brownish  yellow 

Red 
Colorless 
Gray 
Bottle-green 

Copper 
Manganese 
Nickel 
Iron 

SYSTEMATIC  ANALYSIS 


125 


If  the  substance  is  heated  in  the  oxidizing  flame  with  the 
microcosmic  bead  and  floats  on  the  bead  as  an  undissolved 
skeleton,  it  indicates  silicon  dioxide  or  silicates. 

Several  of  the  rarer  elements  also  give  colored  beads.     (See 

Part  V,  p.  139-) 

The  Flame  Test 

A  little  of  the  substance  is  moistened  with  concentrated  sul- 
phuric acid  and  introduced  into  the  lower  edge  of  a  Bunsen  flame 
by  means  of  a  platinum  wire.  If  the  flame  is  colored  yellow, 
observe  it  through  a  cobalt  glass.  When  the  coloration  of  the 
flame  has  ceased,  moisten  the  wire  with  concentrated  hydro- 
chloric acid  and  heat  again.  This  treatment  results  in  the  vola- 
tilization of  the  sulphates  of  the  alkalies,  while  the  less  volatile 
salts  of  the  alkaline  earths  are  converted  into  the  more  volatile 
chlorides.  The  characteristic  flame  colorations  are  as  follows 
(see  also  Part  V,  p.  139): 


COLOR 


INDICATION 


Yellow  ;  invisible  through  blue  glass. 

Carmine  ;  violet  through  blue  glass. 

Scarlet,  masked  by  barium  flame ;  voilet  through 

blue  glass. 

Yellowish  red ;  greenish  gray  through  blue  glass. 
Yellowish  green. 
Emerald  green. 
Azure  blue. 
Light  blue. 
Violet ;  purple  through  blue  glass. 


Sodium. 
Lithium. 
Strontium. 

Calcium. 

Barium  or  borates. 

Copper. 

Copper  chloride. 

Arsenic  compounds. 

Potassium. 


The  Charcoal  Test 

a.  A  small  portion  of  the  substance  is  heated  on  a  piece  of 
charcoal  before  the  blowpipe.  If  deflagration  takes  place  it  indi- 
cates a  nitrate,  nitrite,  or  chlorate,  or  some  other  substance  rich  in 
oxygen.  If  the  substance  melts  and  runs  into  the  charcoal,  it  indi- 
cates an  alkali  salt.  An  evolution  of  a  garlic  odor  indicates  arsenic. 


126 


QUALITATIVE  ANALYSIS 


b.  The  substance  is  mixed  with  twice  its  own  amount  of  so- 
dium carbonate,  and,  after  moistening  slightly,  is  heated  on  char- 
coal before  the  blowpipe  in  the  reducing  flame.  There  may  be 
obtained  : 


PHENOMENA 


INDICATION 


A.  Metal  without  incrustation  : 

Malleable  button. 
Gray  particles. 

B.  Metal  with  an  incrustation  : 

Brittle  button. 

Malleable  button. 

C.  Incrustation  without  metal : 

White  —  yellow  when  hot. 
White  —  garlic  odor. 
Brown. 

D.  White  infusible  mass  : 

Heated  with  cobalt  nitrate  —  blue. 
Heated  with  cobalt  nitrate  —  pink. 
No  color  change. 

E.  Sulphur  compounds   are    reduced   to 

sulphides.      The    melted    mass    is 
scraped  off  and  placed  on  a  silver 
coin,  with  a  drop  of  water. 
A  brown  stain. 


Gold,  silver,  tin,  copper,  lead. 
Iron,  nickel,  cobalt. 

Antimony — with  white  sublimate. 
Bismuth — with  yellow  sublimate. 
Lead  —  with  yellow  sublimate. 

Zinc. 

Arsenic. 

Cadmium. 

Aluminium. 
Magnesium. 
Calcium,  strontium,  barium. 


Sulphur  compounds. 


(For  a  very  full  discussion  of  the  use  of  the  blowpipe  on  charcoal,  see  Moses 
and  Parsons's  Mineralogy,  p.  97  et  seq.} 

These  preliminary  reactions  will  give  enough  information 
concerning  the  character  of  the  sample  to  enable  the  analyst  to 
proceed  with  the  preparation  of  the  solution  for  the  wet  analysis, 
if  such  is  necessary. 

Preparation  of  the  Sample 

We  distinguish  three  cases : 

i.    The  sample  is  a  liquid,  i.e.  a  solution. 


SYSTEMATIC  ANALYSIS 


127 


Evaporate  a  portion  to  dryness  on  a  water  bath.  If  a  residue 
is  obtained,  test  the  solution  for  nitric  acid  and  for  organic 
matter.  If  found,  evaporate  to  dryness  and  ignite  and  redis- 
solve  in  water,  with  addition  of  acid  if  necessary.  Analyze  ac- 
cording to  the  general  method  given  on  page  130.  If  no  residue 
is  found,  test  need  only  be  made  for  acids,  ammonia,  and  hydro- 
gen peroxide. 

2.    The  sample  is  an  alloy  or  metal. 

Break  the  sample  into  small  pieces  or  roll  it  into  thin  sheets. 
Take  a  one-gram  portion  and  treat  with  nitric  acid  (one  part  con- 
centrated HNO3  to  one  part  water). 

A.  No  apparent  action  takes  place.     Treat  the  sample  with 
aqua  regia  and  examine  for  the  noble  metals  and  aluminium. 

B.  Complete  solution  takes  place.     Evaporate  to  near  dryness 
and  dilute  with  water  and  analyze  as  in  the  group  separations 
for  metals,  p.  1 30  et  seq.     The  only  acids  which  need  be  looked 
for  are  silicic,  phosphoric,  and  sulphuric. 

C.  Complete  solution  does  not  take  place.     Evaporate  to  dry- 
ness  as  in  B ;  dilute  with  water,  filter,  and  separate  as  in  the  fol- 
lowing table  : 


RESIDUE 


FILTRATE 


SiO2,  SnO2,  Sb2O5,  P2O5,  Bi2O3, 

traces  of  Pb,  Cu,  Fe,  etc. 
Wash  with   water,  treat  with  am- 
monium sulphide,  boil,  and  filter. 


Analyze  as  in  B  above. 


RESIDUE 


FILTRATE 


SiO2,  Bi2S3,  PbS  CuS 
Dissolve  in  HC1  and  filter.     Add 
filtrate  to  ammonium  sulphide  solu- 
tion and  examine  residue  for  silicic 
acid. 


(NH4)3SbS4,  (NH4)3P04, 

(NH4)2SnS3. 

Treat  with  HC1  and  filter.  Exam- 
ine filtrate  for  phosphoric  acid.  Ex- 
amine residue  for  Group  II  (metals). 


128  QUALITATIVE  ANALYSIS 

3.    The  substance  is  nonmetallic. 

A.    Solution  for  metal  analysis. 

Treat  very  small  portions  of  the  finely  powdered  and  thor- 
oughly mixed  substance  in  a  test  tube,  successively  with  water, 
dilute  hydrochloric  acid,  concentrated  hydrochloric  acid,  dilute 
nitric  acid,  concentrated  nitric  acid,  and  aqua  regia  until  solu- 
tion is  obtained.  If  complete  solution  is  obtained,  a  larger  quan- 
tity —  about  one  fourth  gram  —  of  the  substance  is  dissolved  in 
the  solvent  found  necessary,  and,  if  nitric  acid  or  aqua  regia 
was  used,  is  evaporated  to  almost  complete  dryness  ;  dilute  with 
water  and  analyze  as  directed  in  the  general  analysis  of  the 
groups  on  page  1 30  et  seq. 

If  the  treatment  above  outlined  is  not  sufficient  to  effect  solu- 
tion, the  residue  must  be  fused.1  In  the  absence  of  the  sub- 
stances described  in  a,  b,  c,  etc.,  below,  fusion  may  be  performed 
as  follows :  Take  about  one  half  gram  of  the  residue  insoluble 
in  acids  and  four  to  five  times  the  quantity  of  sodium  carbonate 
and  fuse  on  platinum  until  a  quiet  fusion  is  obtained.  Then 
boil  with  water  and  filter.  The  aqueous  solution  may  be  acidi- 
fied and  added  to  the  solution  obtained  by  acids  if  no  precipitate 
is  formed;  otherwise  it  must  be  analyzed  separately.  The 
residue  insoluble  in  water  is  washed  and  treated  with  strong 
hydrochloric  acid  and  this  solution  added  to  the  one  obtained 
by  acids  if  no  precipitate  is  -formed ;  otherwise  it  is  analyzed 
separately. 

a.  Ignition  with  free  access  of  air  will  dispose  of  sulphur 
and  carbon.  A  test  may  be  made  by  heating  in  a  test  tube 
with  CuO  and  passing  the  gas  evolved  through  very  dilute 
acidified  permanganate  and  into  limewater. 

iThe  ordinary  substances  insoluble  in  water  and  acids  which  may  be  found  in 
the  residue  for  fusion  are  C,  S,  BaSO4,  SrSO4,  CaSO4,  PbSO4,  AgCl,  CaF2,  SiO2,  natural 
and  ignited  oxides  of  Al,  Cr,  Fe,  and  Sn,  many  silicates,  silicon,  and  silicides.  The  pre- 
liminary examination  will  have  indicated  which  of  these  may  be  expected  to  be  present, 
and  the  treatment  of  the  insoluble  residue  must  be  varied  accordingly. 


SYSTEMATIC  ANALYSIS  129 

b.  Silver  chloride  will  be  dissolved  by  a  solution  of  potas- 
sium cyanide  and  reprecipitated  by  ammonium  sulphide. 

c.  Chromic  iron  and  chromic  oxide  are  but  little  affected  by 
fusion  with  sodium  carbonate,  and  may  be  brought  into  solution 
by  fusion  with  an  oxidizing  flux,  such  as  sodium  carbonate  and 
potassium   nitrate,  or  sodium  peroxide.     In  the  latter  case  a 
platinum  vessel  must  not  be  used,  but  the  fusion  is  to  be  carried 
out  in  a  nickel,  iron,  copper,  or  silver  vessel. 

d.  In  case  of  alumina  and  difficultly  soluble  iron  oxides,  the 
best  flux  is  acid  potassium  sulphate. 

e.  In  case  of  calcium  fluoride,  it  is  best  to  treat  with  sul- 
phuric acid  and  then  deal  with  the  sulphate. 

/.  In  case  of  stannic  oxide  or  of  silicon  or  silicides,  it  is 
best  to  fuse  in  a  silver  vessel  with  potassium  hydroxide  and 
then  dissolve  in  water. 

B.    Solution  for  acid  analysis. 

It  is  impossible  to  prepare  a  solution  suitable  for  all  the  acid 
tests.  Group  I  tests  are  all  made  upon  separate  small  portions 
of  the  solid  substance.  A  solution  of  the  remaining  acids  may 
be  prepared  by  boiling  the  sample  for  fifteen  minutes  with  a 
saturated  solution  of  sodium  carbonate  and  filtering  as  described 
in  Part  III,  p.  115.  The  filtrate  will  contain  the  acid  radicals 
of  all  the  common  acids,  except  where  the  original  substance 
is  barium  sulphate  and  certain  silicates. 

The  group  reagent  tests  are  then  made  upon  this  solution, 
and  each  group  found  represented  is  analyzed  as  per  directions 
in  Part  III. 


GENERAL  ANALYSIS  OF  THE  METALS 
Group   I 

To  the  neutral  or  acid  solution  add  dilute  hydrochloric  acid, 
with  constant  stirring,  until  a  precipitate  is  no  longer  formed. 
Filter  and  wash  twice  with  cold  water.  The  nitrate  may  con- 
tain members  of  Groups  II  to  V.  The  precipitate  is  analyzed 
for  metals  of  Group  I  as  indicated  below : 

The  precipitate  is  washed  with  boiling  water  until  no  more  dissolves. 


Precipitate:  AgCl,  HgCl 


Filtrate:  PbCL 


NH4OH. 

Divide  into  several  portions. 
(#)    Treat  with  H2SC>4  ;  a  white 

Residue:  Hg  and 

Filtrate  : 

precipitate  indicates  lead. 

HgNH.Cl 

Ag(NH^)2Cl 

(£)    Pass  in  H2S  ;  a  blaclc  pre- 

A   black  "   residue 

Acidify       with 

cipitate  indicates  lead. 

shows  the  pres- 

HNOr    A    white 

(c}    Treat     with     K2CrO4;     a 

ence  of  mercury 

precipitate     shows 

yellow  precipitate  indicates  lead. 

(ous). 

the  presence  of  Ag. 

Group  II 

The  nitrate  from  Group  I  is  evaporated  nearly  to  dryness,  with 
frequent  additions  of  small  quantities  of  HC1  to  expel  any 
HNO3  which  may  be  present.  Dilute  with  water  and  add  HC1 
to  about  one  part  in  ten,  heat  to  boiling,  and  pass  in  H2S  slowly 
as  long  as  a  precipitate  forms.  *  Filter.  Cool,  dilute,  and  pass 

1  If  arsenates  are  present,  acidify  strongly  and  pass  H2S  ;  filter  and  evaporate  filtrate 
nearly  to  dryness,  dilute,  and  again  pass  H2S  until  the  precipitation  is  complete. 

130 


S 


GENERAL  ANALYSIS  OF  THE  METALS 


133 


o-a 


III 
|ii 

^It 


Is    g 


{I 


g-2 


o  S      -2  *0  £ 

^-l*lall   IM  = 

•3  gf"S    c?    -3     Erf0 

'0*«00*l|f5f    a8- 

!§l4i!sll  lgl 

'n:^oE-3^=S    OX:T) 

'  IsK.--^  els 

lfes!8ili^ 

^«£^<:S    i^:^ 


color  shows  the 
f  manganese. 


_•  v  o 

u-s.s 


H-'IsJ 


13  a.2 


»  - 


S"S 

fiii! 

U  •'S  O  C 

•-  '§*  'i 

J3   W  1)   3 

^  a£"5 


^.s  —     sj  -o «   -     «-y<« 
g  •§  E  I  §  §   .  ^fi  2  d 


fill 


be 


Residue 
iS  (bla 
oS  (blac 


C 
S 


Tes  a  bora 
Bro  ad  shows  t 
presence  of  nickel. 
Blue  bead  shows  th 
presence  of  cobalt.  N  icke 
may  be  present. 


wit 
n  b 


^  dp  itate  :  Ni  (  OH 

Test  with  a  borax  be 
Brown  bead  shows  t 
esence  of  nickel. 


JTsCo(CN} 
id  with  HCI 
o  a  small  bul 
h  a  borax  bea 
icates  Co. 


te 
wi 
n 


Filtrate 
Make 
vaporat 
nd  test 
Blue  i 


134  QUALITATIVE  ANALYSIS 

Whether  a  precipitate  forms  or  not,  add  to  the  solution  while 
still  warm  (without  filtering)  (NH4)2S.  Filter  at  once  and  wash 
the  precipitate  with  water  containing  a  little  (NH4)2S.  The 
filtrate l  may  contain  metals  of  Group  IV  and  V.  The  pre- 
cipitate contains  all  the  metals  of  Group  III  which  are  present. 
(If  phosphoric,  silicic,  boric,  hydrofluoric,  or  tartaric  acids  are 
known  to  be  present,  use  the  following  scheme  for  analysis.) 

A.  If  tartaric  or  oxalic  acid  is  present,  evaporate  the  solu- 
tion or  filtrate  from  Group  II  to  dryness  and  fuse  with  Na2CO3 
and  KNO3,  dissolve  the  fused  mass  in  water,  acidify  with  HC1, 
pass  in  H2S  to  reduce  any  chromic  acid  if  present  to  a  chromium 
salt,  and  treat  as  given  below,  B. 

B.  If  tartaric  and  oxalic  acid  are  absent,  boil  the  solution  or 
filtrate  from  Group  II  to  expel  any  H2S  present.     Add  a  few 
drops  of  HNO3  and  boil.     If  the  solution  becomes  yellow,  iron 
is    probably   present.     Add   a   small   quantity  of    NH4C1  and 
NH4OH  to  alkaline  reaction  and  boil.     If  a  precipitate  forms, 
one  or  more  of  the  metals  Fe,  Cr,  or  Al  are  present.     If  no 
precipitate  forms,  these  are  absent,  and  subsequent  tests  need 
not  be  applied.     In  either  case  add  to  the  solution  while  still 
warm  (without  filtering)  (NH4)2S.     Filter  and  wash  with  water 
containing  a  little  (NH4)2S.     The  precipitate  is  to  be  analyzed 
for   Group  III,   p.  I35.2     The    filtrate  may  contain   metals  of 
Groups  IV  and  V. 

The  precipitate  may  contain  hydroxides  of  Al,  Cr,  and  the 
sulphides  of  Fe,  Co,  Ni,  Mn,  and  Zn,  and  phosphates  and 
similar  compounds  of  Ba,  Ca,  Sr,  &$.d  Mg. 


1  If  the  filtrate  is  dark  brown  in  color,  due  to  some  NiS  remaining  in  solution,  the 
presence  of  Ni  is  indicated.     If  this  is  the  case,  the  filtrate  should  be  acidified  with  acetic 
acid  and  boiled  for  some  time  —  filtered  on  a  separate  filter  and  tested  directly  for  Ni  with 
a  borax  bead.    The  filtrate  contains  members  of  Groups   IV  and  V.     See  Ostwald, 
Scientific  Foundations,  p.  24. 

2  If  phosphoric,  silicic,  boric,  or  hydrofluoric  acids  are  known  to  be  absent,  analyze  by 
the  short  method  on  page  133. 


GENERAL   ANALYSIS   OF  THE  METALS 


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QUALITATIVE  ANALYSIS 


NOTE.  Phosphates,  if  present,  may  be  removed  also  by  evaporating  the  solution  for 
this  group  to  near  dryness,  taking  up  in  concentrated  nitric  acid  and  treating  with  a  strip 
of  tin ;  the  precipitate  is  filtered  off  and  the  filtrate  evaporated  to  dryness  on  the  water 
bath  and  dissolved  in  dilute  hydrochloric  acid.  The  analysis  may  then  be  carried  out 
as  described  on  page  133. 

Group  IV 

The  solution,  or  filtrate  from  Group  III,  is  boiled  until  clear, 
concentrated  to  a  small  bulk,  and  filtered. 

Add  NH4C1,  and  NH4OH  to  alkaline  reaction,  andCNH^COg, 
and  warm  for  about  ten  minutes.  Filter  and  wash  with  hot 
water.  The  precipitate  contains  all  the  members  of  Group  IV 
present.  The  filtrate  may  contain  members  of  Group  V. 


Reserve  a  small  portion  of  the  precipitate  for  flame  tests.  Dissolve  the 
remainder  in  acetic  acid.  Test  a  small  portion  for  Ba  with  K2CrO4.  If  a 
precipitate  forms,  add  K2CrO4  to  the  remainder;  warm,  filter,  and  wash. 


Precipitate  : 


Add  HC1  and 
warm  .  Add 
H2SO4.  If  a  pre- 
cipitate forms, 
Ba  is  present. 
The  precipitate 
will  appear  ye'l- 
low  because  of 
the  K2CrO4  in 
the  solution. 


Filtrate  :     Ca  (  C2 


Sr(CiHz  <92) 


Render  alkaline  with  NH4OH  and  add  (NH4)2CO3. 
Precipitate  CaCO3  and  SrCO3.  Filter,  wash  until  white, 
and  discard  filtrate.  Test  for  Sr  by  flame  test. 

Strontium  is  present  :  Dissolve  precipitate  in  acetic 
acid.  Add  dilute  solution  of  (NH4)2SO4,1  warm,  and  stir 
for  a  few  minutes.  Filter  off  precipitate  and  test  filtrate 
for  Ca  by  adding  NH4OH  and  (NH4)2C2O4.  If  a  pre- 
cipitate forms,  it  shows  the  presence  of  Ca. 

Strontium  is  absent  :  Dissolve  precipitate  in  acetic 
acid,  [make  alkaline  with  NH4OH,  and  add  (NH4)2C2O4. 
A  white  precipitate  shows  presence  of  Ca. 


Group  V 

A.  To  a  portion  of  the  solution,  or  filtrate  from  Group  IV, 
concentrated  to  a  small  volume,  add  NH4OH  in  excess,  NH4C1, 

!If  to  a  portion  of  the  solution  a  saturated  solution  of  CaSO4  is  added  instead  of 
(NH4)2SO4,  a  precipitate  will  form  if  strontium  is  present.  The  precipitate  may  then 
be  tested  for  strontium  by  the  flame  test.  This  separation  is  based  upon  the  difference 
in  solubility  of  the  sulphates  of  Ca  and  Sr. 


GENERAL  ANALYSIS  OF  THE  METALS  137 

and  Na^HPO^  Thoroughly  agitate,  scratch  the  walls  of  the 
vessel  with  a  glass  rod,  and  allow  to  stand  for  some  time.  A 
white  precipitate  shows  the  presence  of  magnesium. 

B.  The  remainder  of  the  solution,  or  filtrate  from  Group  IV, 
is  made  acid  with  H2SO4  and  evaporated  to  dryness  in  a  porce- 
lain dish,  and  heated  at  a  low  red  heat  until  white  fumes  of 
SO3  are  given  off.     Scrape  the  residue  adhering  to  the  sides  of 
the  dish  into  the  center,  where  it  will  be  thoroughly  heated. 
When  cool,  add  water  and  a  drop  of  HC1;  filter  if  not  clear. 
Test  the  filtrate  for  K  with  (a)  Na3Co(NO2)6;    (b)  PtCl4;    (c) 
flame  using  cobalt  glass.     (See  tests  for  K.) 

C.  Test  a  portion  of  the  original  substance  for  sodium  by 
the  flame  test.     A  persistent  yellow  flame  shows  the  presence 
of  sodium. 

D.  Test  a  portion  of  the  original  substance  in  a  test  tube 
by  adding  NaOH  and  warming  gently:     (a)  Odor;     (b)  action 
on  damp  litmus  placed  on  the  mouth  of  the  tube;     (c)  action 
on  rod  moistened  with  HC1  held  in  the  tube.     (See  tests  for 
NH4OH.) 

Acid  Analysis 

No  suitable  table  of  separation  of  the  acid  groups  can  be 
prepared.  The  student  is  referred  to  Group  Analysis  given  in 
Part  III. 


PART  V 

THE    RARE?  METALS 

The  elements  usually  classed  together  as  rare  metals  consist 
in  fact  of  both  metals  and  non-metals.  They  are  of  very  various 
commercial  importance.  Some  are  extremely  widespread  in 
nature,  but  occur  in  very  small  quantities,  e.g.  gold.  Others 
occur  in  considerable  masses,  but  are  not  of  wide  distribu- 
tion, e.g.  tungsten.  Some  of  these  elements  are  of  such  ex- 
tensive use  as  to  be  very  familiar,  e.g.  cerium  and  thorium  in 
the  Welsbach  burner  and  gold  and  platinum  in  their  varied  appli- 
cations, and  others  are  very  rarely  encountered  even  in  the  most 
thoroughgoing  analytical  work,  e.g.  radium.  The  development 
of  industrial  uses  sometimes  makes  'an  otherwise  unfamiliar  sub- 
stance of  considerable  interest  to  the  analyst.  An  example  of 
this  is  furnished  by  the  recent  use  of  tungsten  and  of  tantalum 
as  filaments  in  electric  light  bulbs.  The  study  of  these  elements 
is  therefore  of  continually  increasing  importance  from  the  tech- 
nical veiwpoint,  and  no  analytical  course  is  complete  without 
providing  some  experience  in  detecting  the  presence  of  these 
elements  and  in  the  methods  of  identifying  them.  The  purpose  of 
this  discussion  is  to  present  a  brief  outline  of  the  procedure 
and  the  more  striking  of  the  analytical  reactions  to  the  end  that 
the  student  may,  at  the  worst,  be  able  to  look  intelligently  for 
more  detailed  information  in  parger  works.  In  the  preliminary 
reactions,  the  more  commonly  associated  elements  will  also  be 
given  as  a  kind  of  warning  as  to  when  given  elements  may  be  ex- 
pected to  be  present  in  an  unknown  substance. 

Note.  Fuller  discussions  of  these  elements  will  be  found  in  Treadwell  and  Hall,  Bottger 
Browning's  Introduction  to  the  Rarer  Metals  and  in  Abegg's  Handbuch. 

138 


THE  RARE  METALS  139 

The  rarer  elements  may  be  divided  into  five  groups  as  given 
below.  While  this  division  is  on  the  basis  of  the  group  separa- 
tions of  the  commoner  elements,  it  must  be  remembered  that  it  is 
not  so  sharp  nor  so  satisfactory.  Several  of  the  elements  are 
precipitated  by  the  group  reagents,  not  as  cations,  but  as  acid 
anhydrides,  and  in  other  cases  the  precipitation  of  a  given  element 
may  occur  in  more  than  one  group,  according  to  conditions  or  the 
form  of  compound  of  the  element  undergoing  examination.  This 
grouping  is  chiefly  valuable  as  a  means  of  determining  what  rare 
elements  are  to  be  more  particularly  sought  when  phenomena 
not  accounted  for  in  the  regular  examination  are  encountered. 

The  following  tables  of  flame  and  bead  colorations,  taken  from 
Bottger,  will  be  found  useful  in  preliminary  examination.  Ref- 
erence may  also  be  made  to  Moses  and  Parsons,  pp.  92-94. 

Flame  Test1 

Green  Thallium 

Blue  green  Tellurium  (burning  element) 

Blue  (cornflower)          Selenium,  odor  of  horse  radish 

(burning  element) 
Violet  blue  Indium 

Violet  Rubidium 

Carmine  Lithium 

(See  also  Moses  and  Parsons,  p.  93.) 

Microcosmic  Bead  Tests1 
(See  also  Part  IV,  p.  124.) 

ELEMENT  OXIDIZED  BEAD  REDUCED  BEAD 

Molybdenum  colorless  green  (grass  to  bluish) 

Uranium  yellow  green  (fluor-  green 

escent) 

Vanadium  yellow  green 

Cerium  pale  yellow,  colorless 

hot,  orange  yellow 
Tungsten  colorless  blue  (brownish  red 

with  iron) 

Titanium  colorless  violet  (brownish  red 

with  iron) 

i  O.  Lutz,  Uber  die  Analytische  Verwendung  gefarbter  Glaser,  Zeit.  f.  Analytische 
Chtm.,  Vol.  47,  pp.  1-36 ;  1908. 


I40  QUALITATIVE  ANALYSIS 

These  elements  may  be  grouped  as  follows  : 

GROUP  I.  The  Hydrochloric  Acid  Group :  Thallium,  Molyb- 
denum, Tungsten,  Tantalum,  and  Niobium  (Columbium). 

GROUP  II.  The  Hydrogen  Sulphide  Group  :  Gold,  Platinum, 
Palladium,  Osmium,  Rhodium,  Ruthenium,  Iridium,  Sele- 
nium, Tellurium,  Germanium,  and  (Molybdenum). 

GROUP  III.  The  Ammonium  Sulphide  Group :  Titanium, 
Zirconium,  Uranium,  Indium,  Gallium,  Vanadium,  Beryllium, 
Thorium,  Cerium,  Lanthanum,  Didymium  (Neo-  and  Praeso-), 
Yttrium,  Ytterbium,  Scandium,  and  Erbium. 

GROUP  IV.     The  Ammonium  Carbonate  Group  :  Radium. 

GROUP  V.     The  Soluble  Group  :  Lithium,  Rubidium,  Caesium. 

Group  I 

77,  Mo,  W,  Ta,  Nb  (Cb} 

GENERAL  STATEMENT.  Of  these  elements  only  thal- 
lium is  precipitated  as  a  chloride.  The  others  are  precipitated 
as  acid  anhydrides  from  solutions  in  which  their  salts  are  present. 
These  anhydrides  are  in  most  cases  somewhat  amphoteric  and 
hence  go  into  solution  to  a  large  degree  in  excess  of  acid.  There- 
fore they  may  also  appear  in  other  groups. 

Thallium.  Thallium  is  very  rare,  yet  occurs  in  minute  quan- 
tities in  many  sulphides  and  arsenic  minerals.  It  may  be  found 
in  flue  dust  formed  by  the  roasting  of  pyrites  in  the  manufacture 
of  sulphuric  acid  and  in  metallurgical  operations.  It  may  be 
found  in  optical  glasses  or  artificial  stones  where  it  is  sometimes 
used  on  account  of  the  high  refractive  index  which  it  imparts  to 
the  glass.  It  forms  two  series  of  salts,  univalent  and  trivalent. 
The  trivalent  salts  are  unstable  and  furnish  reactions  resembling 
those  of  aluminium.  The  univalent  compounds  closely  resemble 
those  of  silver.  The  instability  of  thallic  salts  is  such  that  in 
boiling  aqueous  solution  they  are  ordinarily  decomposed  into 


THE  RARE  METALS  14 1 

either  thallic  hydroxide  or  thallbus  salts.  If  'flue  dust,'  or  thal- 
lium glass  after  fusion  with  sodium  carbonate,  is  treated  with 
sulphuric  acid,  the  nitrate  will  give  the  following  characteristic 
reactions  if  thallous  sulphate,  T12SO4  is  present : 

1.  HC1,  HBr,  HI,  or  their  salts  precipitate  thallium  halides 

insoluble  in  ammonium  hydroxide,  acids,  or  potas- 
sium cyanide,  but  soluble  in  sodium  thiosulphate. 
The  halides  are  sensitive  to  light.  The  least  soluble 
is  the  iodide.  The  chloride  is  soluble  in  hot  water 
along  with  lead  chloride,  and  after  the  latter  is  re- 
moved by  dilute  H2SO4,  the  nitrate  gives  a  yellow 
precipitate  with  KI. 

2.  Flame  Test.     On   a  platinum  wire  the  thallium  salts 

color  the  flame  green.  The  test  may  be  made  on 
the  mixed  chlorides  of  Group  I. 

3.  H2S,  in  neutral  solution,  or  soluble  sulphides  precipitate 

Tl^S;  soluble  in  mineral  acids,  but  not  in  acetic 
acid. 

4.  Alkali  hydroxides  cause  no  precipitation,  and  soluble 

carbonates  do  so  only  in  concentrated  solution. 

Molybdenum.  Molybdenum  is  found  rather  frequently  as 
the  sulphide,  molybdenite ;  or  as  lead  molybdate,  wulfenite.  The 
sulphide  closely  resembles  graphite  in  appearance.  Molybde- 
num is  used  in  the  manufacture  of  steel,  both  with  and  without 
nickel,  and  renders  the  steel  hard  and  tough  and  hence  suitable 
for  the  manufacture  of  propeller  shafts,  heavy  guns,  shells,  etc. 
It  is  also  used  for  wire.  Ammonium  molybdate  and  molybdic 
anhydride  are  important  laboratory  reagents.  The  sodium  salt 
is  used  for  blue  glazes  in  pottery  manufacture. 

Molybdenum  forms  the  oxides  MoO,  Mo2O3,  MoO2,  and  MoO3. 
Of  these  the  first  three  are  essentially  basic  and  the  fourth  is  an 
acid  anhydride  and  is  the  most  important,  since  it  is  formed  by 


142  QUALITATIVE  ANALYSIS 

the  use  of  strong  oxidizing  agents  upon  the  element  or  upon 
most  of  its  compounds.  The  anhydride  is  soluble  in  both 
acids  and  alkalies.  Using  a  solution  of  ammonium  molybdate, 
the  following  reactions  are  characteristic : 

1.  HC1,  or  other  dilute   acids,  precipitate  white  H2MoO4 

soluble  in  excess  of  acid  or  alkali.    (See  Tungsten.) 

2.  H2S  added  to  acid  solutions  produce  first  a  blue  colora- 

tion of  the  solution,  and  finally  a  brown  precipitate 
of  both  MoS2  and  MoS3;  soluble  in  ammonium 
sulphide,  forming  sulphomolybdate,  reprecipitated 
by  dilute  acids. 

Since  the  sulphide  is  somewhat  soluble,  it  may 
pass  through  to  Group  IV,  and  when  the  filtrate 
from  the  third  group  is  acidified,  the  brown  MoS3 
will  be  precipitated.  The  sulphide  is  soluble  in 
nitric  or  sulphuric  acids,  forming  molybdic  acid. 

3.  Na2HPO4  precipitates  in  strongly  acid  solutions  yellow 

phosphomolybdate.  See  Reaction,  4,  p.  113;  also 
4,  p.  114. 

4.  H2SO4  evaporated  to  dryness  with  a  trace  of  molybde- 

num compounds  produces  an  intense  blue  coloration. 

5.  On  adding  zinc,  SnCl2,  and  other  reducing  agents,  and 

acidifying  the  solution  with  HC1  or  H2SO4,  a  blue 
coloration  turning  green  and  finally  brown  is  pro- 
duced. 

6.  Bead  test.     The  borax   bead  is  dark  brown  in  the  re- 

ducing flame  and  colorless  in  oxidizing  flame.  See 
p.  124;  also  p.  139. 

Tungsten.  The  tungsten  minerals  of  importance  are  Wolfram- 
ite (Fe,  Mn)WO4,  Scheelite,  CaWO4,  and  Hiibnerite,  MnWO4. 
It  is  used  extensively  in  the  manufacture  of  the  '  self-hardening  ' 
tool  steels,  which  do  not  lose  their  temper  when  hot,  in  the  manu- 


THE  RARE  METALS  143 

facture  of  magnets,  with  aluminium  as  an  alloy  having  high  ten- 
sile strength  and  low  specific  gravity,  for  the  filaments  in  the 
tungsten  lamp  and  in  bronzes  which  have  very  variable  pur- 
poses. Sodium  tungstate  is  used  in  fireproofing  muslin  and 
other  inflammable  materials,  and  as  a  mordant  in  dyeing.  Lead 
tungstate  may  be  substituted  for  white  lead  as  a  pigment.  The 
trioxide  is  a  canary  yellow  pigment. 

Tungsten  forms  two  oxides  WO2  and  WO3,  the  latter  only 
being  of  importance  to  the  analyst.  It  is  an  acid  anhydride,  and 
its  sodium  salt  is  readily  prepared  from  tungsten  minerals  by 
fusion  with  sodium  carbonate  ;  or  N^C^  in  an  iron  vessel.  The 
sodium  tungstate  so  produced  is  soluble  in  water  and  may  be 
used  for  the  following  characteristic  reactions : 

1.  HC1,    HNO3,  and   H2SO4  precipitate    white   hydrated 

tungstic  acid,  becoming  anhydrous  on  boiling  the 
solution  ;  easily  soluble  in  alkalies,  but  only  slightly 
in  strong  acids.  Washed  with  pure  water  it  forms 
a  turbid  colloidal  solution. 

2.  H2S  in   neutral  solution  precipitates   the   sulphide  in- 

completely, soluble  in  (NH^S,  but  not  wholly  re- 
precipitated  by  acids. 

3.  SnCl2  produces  a  yellow  precipitate  which  on  acidification 

with  HC1  and  warming  turns  to  a  beautiful  blue. 
(See  Molybdenum.)  Zinc  and  acid  also  produce  the 
same  blue  precipitate. 

4.  Bead  test.     The  borax  bead  gives  a  brown  color  in  the 

reducing  flame.  The  microcosmic  bead  furnishes  a 
blue  green  bead  in  the  reducing  flame,  colored  blood- 
red  by  a  trace  of  ferrous  sulphate. 

Tantalum  and  Niobium.  Tantalum  is  a  quite  rare  element 
and  usually  occurs  in  conjunction  with  niobium.  The  most  im- 
portant minerals  are  tantalite  and  columbite.  Wolframite  fre- 


144  QUALITATIVE  ANALYSIS 

quently  contains  small  quantities  of  both  tantalic  and  niobic  acid, 
as  does  also  cassiterite  (tin  stone).  The  two  elements  are  quite 
difficult  of  separation,  and  because  of  this  fact  and  their  small 
technical  application  will  be  considered  together.  The  only  ap- 
plication of  tantalum  is  as  a  filament  in  the  tantalum  light  bulb. 
Niobium  has  no  technical  uses. 

NOTE.  The  names  of  the  elements  developed  as  follows :  The  element  niobium  was 
first  discovered  in  a  mineral  found  in  Connecticut  and  was  called  Columbium.  Afterwards, 
tantalum  was  obtained  in  a  mineral  from  Finland  and  called  tantalum  because  of  the  in- 
solubility of  the  oxide,  in  reference  to  the  plight  of  Tantalus  in  Hades.  The  two  elements 
were  supposed  to  be  identical  by  Wollaston  (1809).  In  1844  Rose  separated  an  element 
from  tantalite  which  he  called, niobium  (Niobe,  the  daughter  of  Tantalus),  and  since  this 
element  proved  to  be  identical  with  columbium,  both  names  are  encountered  and  both  Cb 
and  Nb  are  used  as  symbols.  The  name  niobium  is  most  commonly  used. 

The  ores  containing  these  elements  may  be  prepared  for  ex- 
amination as  follows :  Fuse  the  finely  pulverized  mineral  with 
acid  potassium  sulphate  in  a  platinum  crucible  (eight  parts  of  the 
sulphate  to  one  of  mineral).  Boil  the  cooled  mass,  after  pulver- 
izing, with  cold  water  (see  Ti  reactions),  and  repeat  several  times 
with  fresh  quantities.  Filter  and  digest  the  residue  with  ammo- 
nium sulphide  to  remove  tungsten  and  tin,  which  are  very  fre- 
quently associated  elements  in  these  ores.  Boil  with  dilute 
hydrochloric  acid  and  filter  and  fuse  the  residue  in  caustic  potash 
in  a  silver  crucible.  Dissolve  the  fused  mass  with  water  and 
filter.  The  filtrate  will  contain  potassium  niobate  and  tantalate 
and  will  give  the  following  reactions  : 

I.  HC1  precipitates  both  tantalic  and  niobic  acids,  only 
slightly  soluble  in  excess  —  tantalic  more  than 
niobic.  Both  are  soluble  in  concentrated  H2SO4 
on  warming,  and,  on  dilution,  the  tantalic  acid  is 
precipitated  while  the  niobic  is  not. 

2..  (a)  The  tantalic  acid  reprecipitated  from  the  dilute 
sulphuric  acid  dissolves  in  strong  HC1,  from  which 
it  is  reprecipitated  by  ammonium  hydroxide  or  am- 
monium sulphide. 


THE  RARE  METALS  145 

(b)  The  solution  of  hiobic  acid  in  dilute  H2SO4  is  pre- 
cipitated by  ammonium  hydroxide  and  ammonium 
sulphide. 

3.    (a)   Zinc  and   HC1   added  to  the  precipitated  tantalic 
acid  produce  no  color  changes. 

(£)  HF  dissolves  both  the  acids,  if  added  in  excess,  and 
on  addition  of  KF,  the  tantalum  is  precipitated  as 
K2TaF7,  changed  by  boiling  into  the  still  more  in- 
soluble K4Ta4O5F14.  The  double  niobic  fluoride 
and  oxyfluoride  are  very  soluble. 

Group  II 

THE   HYDROGEN    SULPHIDE   GROUP 
Au,  Pt,  Pd>  Os,  Rh,  Ru,  Ir,  Se,  Te,  Ge,  and  (Mo) 

The  elements  listed  in  this  group  are  all  precipitated  by 
H2S  in  acid  solution.  The  first  two  are  sometimes  dis- 
cussed with  the  common  metals  because  of  their  frequent  use 
in  the  laboratory  and  in  the  arts.  They  are  placed  by  us 
among  the  rare  elements  partly  because  of  the  infrequent  use 
of  the  ordinary  qualitative  methods  for  their  detection,  fire 
assay  methods  being  usually  employed,  and  also  because  of 
the  association  in  nature  of  platinum  with  several  other  ele- 
ments of  the  group.  Selenium  and  tellurium  properly  are 
nonmetals,  but  their  place  here  is  secured  by  the  formation 
of  sulphides.  Detailed  methods  of  se'paration  will  not  be 
given,  but  the  reader  is  referred  to  Bottger's  Qualitative 
Analyse,  2d  Auflage,  p.  473,  for  the  separation  of  the  platinum 
metals,  or  to  Mylius  and  Dietz,  Ber.  d.  deutsch.  Chem.  Ges., 
Vol.  31,  p.  3187,  and  to  Bottger,  p.  480,  for  the  remainder  of 
the  group,  or  Noyes  and  Bray,  Jour.  Am.  Chem.  Soc.,  Vol.  29, 
pp.  137-205. 

Gold.  Gold  is  usually  found  native,  though  in  small  amounts, 
disseminated  through  quartz,  sand,  or  other  minerals.  It  is  also 


146  QUALITATIVE  ANALYSIS 

found  associated  with  silver,  as  a  telluride.  All  the  compounds 
are  dissociated  by  heat,  giving  the  free  element.  It  is  liberated 
from  its  salts  through  displacement  by  nearly  all  other  metals. 
It  forms  alloys  readily  with  many  elements  which  are  extremely 
useful  in  the  arts,  especially  in  coinage  and  in  jewelry.  It 
forms  an  amalgam  with  mercury,  which  is  used  as  a  means 
of  collecting  the  gold  from  the  minerals  with  which  it  is  asso- 
ciated. It  also  unites  readily  with  chlorine,  and  forms  a  double 
salt  with  potassium  cyanide,  both  of  which  properties  are  made 
use  of  in  metallurgical  operations. 

The  detection  of  its  presence  in  minerals,  as  well  as  the 
estimation  of  its  amount,  is  ordinarily  performed  by  the  fire 
assay  method,  for  which  consult  works  on  assaying.  It  forms 
two  series  of  salts,  behaving  as  a  univalent  and  as  a  trivalent 
metal.  It  is  not  attacked  by  acids,  but  if  treated  with  aqua 
regia  it  dissolves,  forming  auric  chloride,  AuCl3,  soluble  in  water. 
The  solution  may  be  used  for  the  following  reactions  : 

1.  KOH  or  NaOH  precipitates,  if  added  slowly,  reddish 

brown  Au(OH)3,  soluble  in  excess,  forming  an 
aurate. 

2.  NH4OH  precipitates  fulminating  gold,  A^NH^OH^ 

or  AuN2H3  •  3H2O. 

3.  H2S  precipitates,  in  the  cold,  black  gold  sulphide,  Au2S2, 

soluble  in  (NH4)2S4  with  some  difficulty;  readily 
soluble  in  K2S,  insoluble  in  acids ;  soluble  in  aqua 
regia.  In  boiling  solutions  metallic  gold  is  precipi- 
tated. 

4.  FeSO4  precipitates  metallic  gold  as  a  brown  powder. 

5.  H2C2O4  precipitates  metallic  gold,  reaction  hastened  by 

warming. 

6.  SnCl2  precipitates,  from  not  too  acid  solutions,  metallic 

gold  as  a  reddish  purple  substance  Purple  of  Cassius). 


THE  RARE  METALS  147 

Platinum.  Platinum  occurs  widely  distributed,  but  in  small 
amounts,  chiefly  in  the  native  state,  but  usually  alloyed  with 
smaller  quantities  of  palladium,  osmium,  iridium,  rhodium,  and 
ruthenium.  It  is  not  largely  produced,  and  by  reason  of  its 
varied  application,  is  continually  increasing  in  value.  On  account 
of  its  small  chemical  activity,  it  is  used  very  extensively  in  labora- 
tories as  wire,  foil,  dishes,  stills,  etc.  It  occludes  gases  readily, 
and  in  the  finely  divided  condition  is  an  excellent  catalytic  agent. 
As  a  laboratory  reagent  its  salts  are  frequently  used,  and  would 
be  used  more  extensively  were  they  less  expensive.  It  finds 
application  in  the  arts  as  the  leading  wires  for  electric  light  bulbs, 
since  its  coefficient  of  expansion  and  that  of  glass  are  nearly 
equal ;  in  the  manufacture  of  standard  electrical  instruments  and 
in  photographic  processes.  It  forms  two  series  of  salts  in  which 
it  acts  as  a  bivalent  and  as  a  tetravalent  metal.  Both  as  a  biva- 
lent and  tetravalent  element,  it  forms  a  part  of  complex  negative 
ions.  The  bivalent  forms,  especially,  are  unstable,  and  it  is  mainly 
with  the  tetravalent  compounds  that  we  have  to  deal.  All  the 
compounds,  like  those  of  gold,  are  decomposed  by  heat,  leaving  the 
metal.  The  element  is  rarely  sought  analytically  in  the  wet  way, 
being  obtained  from  its  ores  by  the  fire  assay  method  when  its 
detection  and  estimation  are  desired.  It  is  insoluble  in  acids,  but 
may  be  obtained  from  associated  gangue  by  digestion  with  aqua 
regia  as  chlorplatinic  acid,  H2PtCl6.  The  solution  may  be  used 
for  the  following  characteristic  reactions  : * 

1.  Potassium  or  ammonium  salts  precipitate  the  difficultly 

soluble  chlorplatinates.     See   Potassium  reactions, 
p.  96.     (Used  to  separate  platinum  from  gold.) 

2.  H2S  precipitates  slowly  in  cold  but  rapidly  in  hot  solu- 

tions, brown  PtS  ;  insoluble  in  acids  ;  soluble  in  aqua 
regia ;  difficultly  soluble  in  the  polysulphides  of  the 

1   For  the  preparation  of  chlorplatinic  acid  in  pure  condition,  see  Treadwell  and  Hall, 
P- 234. 


148  QUALITATIVE  ANALYSIS 

alkali  metals  and  ammonium,  from  which  it  is  re- 
precipitated  on  acidification. 

3.    FeSO4  H2C2O4,  and  SnCl2  do  not  precipitate  metallic 
platinum. 

The  platinum  metals,  Pd,  Os,  Rh,  Ru,  and  Ir,  occur  practically 
always  as  constituents  of  the  crude  platinum  found  in  nature 
and  are  separated  from  it  in  its  purification.  Their  separation 
from  each  other  is  a  difficult  analytical  operation,  and  the  student 
has  already  been  referred  elsewhere  for  details.  See  p.  145. 
The  present  discussion  will  therefore  be  very  brief  and  confine 
itself  to  the  mode  of  obtaining  a  solution  of  each  and  to  a  very 
few  of  the  most  striking  characteristics. 

Palladium.  Palladium  is  sometimes  alloyed  with  gold  and 
silver  as  an  ore,  but  is  usually  associated  with  platinum.  It  is 
sometimes  used  to  make  graduated  scales  for  astronomical  in- 
struments, to  plate  silverware,  and  in  dentistry.  It  is  used  also 
to  manufacture  automatic  gas  lighters,  because  of  its  power 
of  occlusion,  which  property  also  makes  it  useful  for  gas 
analysis. 

While  palladium  forms  two  series  of  compounds  in  which  it 
is  bivalent  and  tetravalent,  only  the  former  are  sufficiently  stable 
for  analytical  purposes.  It  is  the  only  metal  of  the  group  which 
dissolves  in  nitric  acid,  and  the  aqueous  solution  of  palladous 
nitrate  presents  the  following  reactions  : 

1.  H2S  precipitates  black  PdS  from  acid  and  neutral  solu- 

tions, insoluble  in  ammonium  sulphide,  but  soluble 
in  boiling  concentrated  HC1. 

2.  KI  produces  a  black  precipitate,  PdI2,  soluble  in  excess 

of  precipitate  and  also  in  ammonium  hydroxide,  a 
characteristic  reaction. 

3.  Hg(CN)2  precipitates  Pd(CN)2,  insoluble  in  HC1,  but 

soluble  in  KCN  and  NH4OH. 


THE  RARE  METALS 


149 


Osmium.  Osmium  finds  some  technical  application  as  a  lamp 
filament,  and  its  tetroxide  is  used  in  bacteriological  operations. 
The  element  manifests  a  valency  of  2,  3,  4,  6,  and  8,  the  last 
being  the  best  known  in  the  form  of  the  oxide  OsO4,  which  is 
known  as  osmic  acid,  though  it  has  scarcely  any  acid  character. 
When  platinum  alloys  are  treated  with  aqua  regia,  osmium  and 
iridium  remain  undissolved,  and  the  osmium  may  be  separated 
by  volatilization  in  a  current  of  oxygen.  If  in  a  finely  divided 
condition,  it  is  soluble  in  aqua  regia  and  can  be  driven  out  of 
the  solution  by  distillation.  It  is  characterized  by  an  offensive 
chlorinelike  odor  and  irritating  reaction  with  the  mucous  mem- 
brane of  eyes  and  throat.  The  solution  of  osmic  acid  in  KOH 
may  be  used  for  the  following  reactions : 

1.  Treated  with  nitric  acid,  the  osmic  acid  is  liberated,  and 

its  odor  may  be  noted. 

2.  H2S  precipitates  in  HC1  solution  the  sulphide,  insoluble 

in  (NH^S. 

3.  Indigo  is  decolorized  by  osmic  acid. 

4.  SnCl2  precipitates  a  brown  precipitate  soluble  in  HC1  to 

a  brown  solution. 

Iridium.  Osmium  iridium  alloy  is  used  because  of  its  hard- 
ness, infusibility,  and  indifference  to  reagents,  as  pen  points, 
watch  and  compass  bearings,  and  iridium  itself  may  be  used  for 
knife  edges,  for  balances,  and  similar  purposes.  Iridium  remains 
nearly  insoluble  when  the  platinum  alloys  are  digested  in  aqua 
regia.  The  residue  not  volatile  in  a  current  of  oxygen  can  be 
iridium  and  ruthenium,  and  may  be  fused  in  a  silver  crucible  with 
sodium  hydroxide,  and  the  fused  mass,  after  removal  from  the 
crucible,  may  be  dissolved  in  aqua  regia,  giving  a  solution  of 
lg  which  will  give  the  following  reactions : 

i.    H2S  first  decolorizes  the  solution  and  finally  precipitates 
Ir2S3,  readily  soluble  in  (NH4)2S. 


150  QUALITATIVE  ANALYSIS 

2.  NaOH  changes  the  color  of  the  solution  from  dark  red 

to  green ;  on  warming  it  changes  to  a  sky-blue  color. 

3.  Reducing  agents  usually  change  the  color  to  an  olive- 

green. 

Ruthenium.  Neither  the  element  nor  its  compounds  find  any 
extensive  use  in  the  arts.  It  is  therefore  only  likely  to  be  en- 
countered in  the  investigation  of  platinum  ores.  It  is  insoluble 
in  aqua  regia. 

Rhodium.  This  metal  likewise  finds  no  special  commercial 
application  and  is  therefore  likely  to  be  of  interest  only  in  the 
analysis  of  platinum  ores  for  which  the  student  is  referred  to 
the  paper  of  Mylius  and  Dietz.  See  p.  145. 

Selenium.  This  element  resembles  sulphur  closely  in  many 
of  its  properties  and  is  therefore  found  replacing  it  in  many  of 
its  compounds.  It  is  of  quite  frequent  occurrence,  but  almost 
invariably  in  very  small  quantities.  It  is  also  sometimes  found 
in  deposits  of  native  sulphur  as  a  sulphide  or  selenide.  On 
roasting  or  burning  it  is  converted  to  the  solid  oxide  SeO2  which 
does  not  readily  oxidize  to  the  higher  form  and  consequently  re- 
mains as  flue  dust  or  solid  deposit  in  the  acid  chambers  in  these 
processes.  As  an  element,  it  finds  application  because  of  the 
effect  of  light  upon  its  electrical  conductivity,  in  the  manufac- 
ture of  electrical  devices.  It  exists  in  two  allotropic  modifica- 
tions and  two  series  of  compounds  analogous  to  those  of  sulphur. 
Selenium-bearing  flue  dust  may  be  fused  with  potassium  hy- 
droxide in  a  silver  crucible  and  extracted  with  water,  and  the 
solution  of  K2SeO3  and  K2Se  used  for  the  following  reactions  : 

i.  H2S  produces  in  dilute  acid  solution  a  lemon-yellow 
precipitate  of  selenium  and  sulphur,  soluble  in  am- 
monium sulphide. 

H2SeO3  +  2  H2S  ->  3  H2O  +  2  S  +  Se. 
The  acidifying   of   the  original  solution   liberates 


THE  RARE  METALS  151 

H2Se,  detected  by  its  characteristic  odor  of  decayed 
cabbage  or  horse-radish. 

i.    Reducing  agents  precipitate  the  selenium  in  the  red 
variety,  changed  by  continued  boiling  to  black. 

3.  Chlorine  or  aqua  regia  converts  the  selenious  acid  to 

selenic  acid,  which  forms  with  BaCl2  insoluble 
BaSeO4.  The  latter  on  boiling  with  HC1  is  re- 
duced, with  evolution  of  chlorine,  to  BaSeO3,  solu- 
ble in  acid. 

4.  Flame  test :    Heated    in  the  reducing  flame,    selenium 

and  its  compounds  color  the  flame  cornflower-blue 
(a  shred  of  asbestos  serves  in  lieu  of  a  wire). 

5.  Codeine  added  to  concentrated  H2SO4  containing  traces 

of  selenium  will  give  a  green  color. 

Tellurium.  Tellurium  is  of  less  frequent  occurrence  than 
selenium,  but  it  is  more  important  because  of  its  occurrence  with 
gold  in  certain  ores.  Since  the  gold  volatilizes.with  the  tellurium 
on  heating,  these  ores  require  special  treatment.  Considerable 
quantities  of  tellurium  are  obtained  at  the  Baltimore  and  Omaha 
works  of  the  American  Smelting  and  Refining  Co.  as  a  by- 
product.1 No  commercial  uses  have  been  found  for  it.  Tellu- 
rium compounds  when  introduced  into  the  human  system  impart 
a  peculiarly  offensive  odor  to  the  breath.  Tellurium  forms  two 
series  of  compounds  analogous  to  those  of  sulphur  and  selenium. 
The  following  reactions  are  characteristic : 

1.  A  fragment  of  tellurium  heated  gently  with  concentrated 

sulphuric  acid  produces  a  beautiful  carmine  color. 

2.  Flame  test :     Tellurium   and  its  compounds,  heated  in 

the  reducing  flame,  give  a  blue  flame. 

3.  HNO3  oxidizes  tellurium  to  tellurous  acid,  which  is  but 

little  soluble  in  water,  but  is  dissolved  readily  by 

1  This  product  probably  contains  two  unidentified  elements.    See  Flint,  Am.  Jour,  of 
Set.,  Vol.  30,  p,  1209;    1910. 


152  QUALITATIVE  ANALYSIS 

bases.  It  is  also  soluble^in  strong  nitric  acid,  form- 
ing tellurium  nitrate,  which  is  hydrolyzed  by  water. 

4.  Tellurium  compounds  heated  with    sodium    carbonate, 

out  of  contact  with  air,  produce  Na2Te ;  soluble  in 
water  to  a  red  solution,  from  which  tellurium  may 
be  precipitated  by  a  current  of  air. 

5.  H2S  and  other  reducing  agents  precipitate  tellurium,  or 

TeS2,  which  is  soluble  in  ammonium  sulphide. 

Germanium.  This  extremely  rare  element  belongs  to  the 
carbon  family  and  in  properties  lies  between  tin  and  carbon. 
Its  chief  interest  lies  in  the  prophecy  of  its  properties  by  Men- 
delejeff  in  1871  and  its  discovery  fourteen  years  later.1  It  is 
precipitated  from  its  soluble  salts  by  hydrogen  sulphide  in 
strongly  acid  solution,  soluble  in  ammonium  sulphide.  The 
sulphide  is  white. 

Group  III 

V,   Ti,  Zr,  Be  (67),  Th,  Ce,  La,  Di  (Ndt  Pr\  Yt,  Yb,  Sc,  Ery 

and  other  still  rarer  elements 

Of  these  elements,  several  are  so  very  rare  as  to  render  their 
discussion  in  these  notes  of  but  little  value,  and  the  student 
is  referred  elsewhere  when  their  detection  is  rendered  neces- 
sary. They  are  all  precipitated  by  ammonium  hydroxide  and 
ammonium  sulphide  along  with  the  other  elements  of  the  third 
group,  although  part  of  them  are  acid-forming  elements  and  are 
precipitated  because  of  the  hydrolysis  of  their  ammonium  salts. 

Vanadium.  This  element  is  found  quite  widely  distributed, 
but  in  conjunction  with  the  compounds  of  other  elements  and 
in  small  amounts.  The  commonest  minerals  are  Vanadinite, 
(Pb6(VO4)3Cl),  Carnotite,  (K2O  2  U2O3  .  V2O5),  Descloizite, 
((PbZn)2(OH)V04). 

1-  Jour,  f.prakt.  Chem.,  Vol.  34,  p.  177. 


THE  RARE  METALS  153 

The  element  is  used  as  an  alloy  in  the  manufacture  of  va- 
nadium steel,  which  has  very  high  tensile  strength.  The  com- 
pounds are  also  used  in  porcelain  manufacture  for  coloring  the 
enamel  green,  in  calico  printing  and  silk  dyeing,  and  in  the 
preparation  of  an  indelible  ink.  The  use  of  vanadium  in  steel 
has  greatly  increased  the  value  of  vanadium  ores. 

The  element  forms  a  series  of  five  oxides,  of  which  the  pen- 
tavalent  form  is  the  most  important.  A  solution  of  vanadium 
suitable  for  examination  may  be  prepared  as  follows :  Fuse 
the  vanadium  mineral  with  KOH  and  extract  with  water.  The 
solution  is  filtered  and  the  second  group  metals  removed  by 
H2S.  The  filtrate  is  boiled  with  a  few  drops  of  nitric  acid  to 
oxidize  the  vanadium  to  vanadic  acid.  The  solution  may  be 
used  as  follows : 

1.  H2S  causes  no  precipitate,  except  sulphur,  but  changes 

the  color  to  blue.    Zinc  produces  the  same  reaction. 

2.  Ammonium  hydroxide  and  ammonium  chloride  slowly 

precipitate  ammonium  metavanadate,  NH4VO3. 

3.  H2O2  added  to  the  nearly  colorless  alkali  vanadates  in 

faintly  acid  solution  produces  a  brownish  red  color. 

4.  Bead  test :  Vanadium  compounds  color  the  bead  brown- 

ish  yellow   in   the   oxidizing  flame,    green   in    the 
reducing  flame. 

Titanium.  Titanium  belongs  in  the  carbon  family  and  is 
fairly  widely  distributed,  both  in  its  own  mineral  forms  as  Rutile, 
Ti(TiO4),  and  Titanic  iron,  FeTiO3,  Titanite,  CaSiTiO3,  etc.,  and 
as  an  impurity  in  many  iron  ores.  It  is  therefore  likely  to  be 
encountered  in  the  analysis  of  iron  ores  or  of  rocks.  The  metal 
itself  finds  but  little  use  except  for  demonstration  of  its  proper- 
ties, and  its  presence  in  other  ores  is  considered  undesirable 
because  of  its  refractory  character,  though  a  titanium  steel  may 
be  made  which  possesses  many  desirable  characteristics.  Its 


154  QUALITATIVE  ANALYSIS 

ores  are  sometimes  used  in  lining  puddling  furnaces,  to  color 
porcelain  yellow,  in  coloring  artificial  teeth,  etc.  The  titanite 
is  used  as  a  gem.  Its  occurrence  and  uses  being  so  relatively 
frequent  and  since  it  is  also  found  in  some  sands,  in  certain 
varieties  of  coal,  in  meteorites,  occasionally  in  ashes,  bones,  etc., 
it  is  sometimes  classed  among  the  common  elements  and  its 
study  and  analytical  reactions  provided  for  in  the  regular  groups. 
It  forms  four  oxides,  of  which  the  dioxide  is  the  most  impor- 
tant and  from  which  its  natural  compounds  are  derived.  Its 
ores  are  not  easily  rendered  soluble  by  fusion  in  sodium  car- 
bonate, and  it  is  best  to  fuse  in  potassium  acid  sulphate  and  dis- 
solve the  fused  mass  in  cold  water.  The  titanium  dissolves  as 
Ti(SO4)2,  and  the  solution  may  be  acidified  with  hydrochloric 
acid  and  used  for  the  following  characteristic  reactions : 

NOTE.  A  preparation  of  pure  titanium  sulphate  is  most  easily  made  as  follows :  The 
fused  mass  is  dissolved  in  cold  water  and  filtered  and  then  heated  to  boiling.  The  pre- 
cipitated TiO(OH)2  is  filtered  hot  and  redissolved  in  concentrated  sulphuric  acid. 

1.  KOH,    NaOH,  NH4OH,  (NH4)2S,  and  carbonates,  in- 

cluding BaCO3,  precipitate  in  the  cold  orthotitanic 
acid,  readily  soluble  in  acids;  hot,  metatitanic  acid 
is  precipitated ;  difficultly  soluble  in  acids. 

2.  NaC2H3O2  forms  Ti(C2H3O2)4,  completely  hydro lyzed 

by  boiling.  (This  hydrolyzing  of  salts  occurs  on 
boiling  the  salts  of  ordinary  minerals  acids,  hence 
the  necessity  of  using  cold  water  in  preparing  the 
solution  above.) 

3.  H2O2  added  to  not  too  strongly  acid  solutions  produces 

a  yellow  to  orange-red  color,  due  to  TiO3.  (See 
also  Vanadium.)  Ferric  chloride,  of  course,  inter- 
feres with  this  reaction,  as  do  also  the  chromates. 

4.  Tin  or  zinc  in  acid  solutions  reduce  the  titanium  to  the 

trivalent  form,  giving  a  violet  color.  Titanium  salts 
are  not  reduced  by  H2S. 


THE  RARE  METALS  155 

5.  Bead  reaction:  Titanium  beads  are  colorless  in  the  oxi- 
dizing flame,  but  violet  in  the  reducing  flame.  By 
addition  of  tin  the  bead  is  reduced  more  quickly. 

Zirconium.  Zirconium  silicate  ZrSiO4  (Zircon)  is  the  com- 
monest natural  form  of  this  element,  though  it  occurs  in  a  number 
of  other  minerals.  The  oxide  of  the  metal  is  unaltered  by  the 
oxyhydrogen  flame  and  glows  brilliantly  when  heated.  It  is  con- 
sequently used  instead  of  lime  in  the  Drummond  light  in  light- 
houses and  is  the  chief  constituent  of  the  glowing  filament  in  the 
Nernst  lamp.  These  uses  are  of  diminishing  importance  because 
of  the  introduction  of  other  methods  of  lighting.  The  mineral 
zircon  is  sometimes  cut  and  used  as  a  gem  under  the  name  hya- 
cinth. Less  valuable  gems  of  the  same  material  are  known  as 
jargon.  The  element  forms  two  oxides  and  corresponding  salts, 
of  which  those  corresponding  to  ZrO2  are  more  important.  A 
convenient  method  of  obtaining  a  solution  for  examination  is  as 
follows:  Fuse  the  pulverized  mineral  substance  with  sodium 
carbonate  at  as  high  a  heat  as  possible.  After  cooling,  digest 
the  mass  with  cold  water  and  boil.  The  sodium  zirconate  is 
hydrolyzed  and  remains  undissolved.  The  residue  may  be  dis- 
solved in  sulphuric  acid  and  after  removal  of  second  group 
metals,  if  present,  the  solution  is  precipitated  with  NH4OH  and 
filtered.  The  precipitate  may  be  redissolved  in  sulphuric  acid 
and  used  for  the  following  tests:  l 

i.  NH4OH,  NaOH,  KOH,  and  (NH^S  precipitate  white 
Zr(OH)4;  insoluble  in  excess ;  soluble  in  acids.  If 
precipitated  hot,  it  is  difficultly  soluble  in  dilute 
but  readily  soluble  in  concentrated  acid. 

l  If  iron  and  alumina  are  also  present,  they  may  be  removed  by  precipitation  with  ex- 
cess of  NH4OH  in  cold  solution  and  filtering.  The  precipitate  dissolved  in  sulphuric 
acid  is  treated  with  an  excess  of  ammonium  carbonate  solution  and  again  filtered.  The 
filtrate  is  boiled,  and  zirconium  hydroxide  is  precipitated. 


156  QUALITATIVE  ANALYSIS 

2.  H2O2  precipitates  white  Zr2O5  from  solutions  not  too 

acid.     The  precipitate  evolves  chlorine  when  treated 
with  HC1.     No  color  is  produced.     (See  Ti.) 

3.  K2SO4  precipitates  from  a  concentrated  not  too  acid 

solution  white  basic  zirconium  sulphate ;    insoluble 
in  dilute  HC1.     (Th  and  Ce  give  similar  reactions.) 

Uranium.  This  element  occurs  in  a  number  of  rare  minerals, 
the  most  widely  distributed  of  which  is  pitchblende,  or  uraninite, 
a  mineral  which  is  75  to  8$  per  cent  uranium  oxide  (or  uranyl 
uranate),  but  contains  also  nitrogen,  helium,  thorium,  zirconium, 
calcium,  iron,  copper,  lead,  etc.  It  is  most  famous  as  the  source 
from  which  the  radioactive  elements  radium  and  polonium  were 
obtained  by  Professor  and  Madame  Curie.  It  is  found  most 
abundantly  at  Joachimstal,  Bohemia,  but  also  occurs  in  many 
places  in  the  United  States.  The  metal  is  used  to  some 
extent  in  the  manufacture  of  gun  steel.  The  oxide  is  used  to 
produce  a  glass  of  a  peculiar  greenish  yellow  fluorescence,  and 
this  furnishes  its  most  extensive  application.  The  salts  are  used 
to  produce  orange  and  black  tints  for  enamels  and  china  paint- 
ing. Some  uranium  also  is  used  in  photography  and  in  the 
preparation  of  the  Welsbach  mantles.  (See  Cerium  and 
Thorium.)  The  element  forms  a  complex  series  of  oxides  and 
corresponding  compounds,  of  which  the  ones  which  chiefly 
concern  the  analyst  are  uranous  salts,  where  the  element  is 
tetravalent  and  the  uranyl,  or  basic,  salts,  where  it  is  hexavalent. 
A  solution  suitable  for  examination  for  uranium  reactions 
may  be  prepared  from  its  ores  by  digesting  the  finely  pulverized 
material  with  aqua  regia,  and  evaporation  to  dryness  on  the 
water  bath.  The  residue  is  extracted  with  water  and  dilute 
hydrochloric  acid  and  treated  with  H2S  to  remove  second- 
group  metals.  The  filtrate  is  boiled  with  nitric  acid  and  filtered. 
The  filtrate  is  treated  with  an  excess  of  ammonium  hydroxide 


THE  RARE  METALS  157 

and  again  filtered.  The  precipitated  ammonium  uranite  will 
dissolve  in  an  excess  of  ammonium  carbonate.  This  solution 
may  be  filtered  and  acidified  with  HC1  and  the  yellow  solution 
used  for  the  following  reactions : 

1.  NH4OH      precipitates      yellow      ammonium     uranate 

(NH^UgOy,  soluble  in  alkali  carbonates. 

2.  K4Fe(CN)6  causes   a   brown   precipitate   or,   in    dilute 

solutions,   a  brown  coloration,   insoluble  in  dilute 
acid. 

3.  H2O2   produces   a   yellowish  white   precipitate  in  con- 

centrated solutions,  if  the  solution  is  kept  neutral 
by  drop-by-drop  addition  of  ammonia. 

4.  Zinc  produces  a  green  coloration. 

5.  Bead  test:     The  uranium  bead  is  yellow  green  in  both 

oxidizing  and  reducing  flames  and  is  fluorescent  in 
the  former. 

Beryllium.  This  element  occurs  rather  widely  distributed, 
but  usually  in  small  quantities,  and  the  specially  fine  crystals 
of  its  compounds  are  cut  and  used  as  gems.  The  most  im- 
portant minerals  are  beryl  (emerald,  aquamarine),  Be3Al2(SiO3)6, 
and  chrysoberyl  (Alexandrite,  cat's  eye),  BeAl2O3;  also  gadol- 
inite,  Be2FeYSi2O10. 

It  is  bivalent  in  all  its  compounds,  and  its  salts  are  sweetish 
in  taste ;  hence  the  name  Glucinum.  It  may  be  obtained  in  so- 
lution by  fusion  of  beryl  with  sodium  carbonate  and  treatment 
of  the  fused  mass  with  hydrochloric  acid.  The  solution  is 
filtered  from  the  silicic  acid  and  evaporated  to  dryness  and  dis- 
solved in  dilute  acid.  This  solution  is  treated  with  an  excess 
of  ammonium  carbonate  and  filtered  cold.  The  solution  is  fil- 
tered and  boiled  and  the  precipitate  of  basic  carbonate  dissolved 
in  HC1.  The  solution  will  not  be  quite  free  from  aluminium, 
but  will  serve  for  the  following  reactions : 


158  QUALITATIVE  ANALYSIS 

1.  NH4OH  and(NH4)2S  precipitate  Be(OH)2  not  soluble 

in   excess,   but  readily  soluble  in  acids   or   alkali 
bases. 

2.  H2C2O4  and  (NH4)2C2O4  produce  no  precipitate  (which 

distinguishes  it  from  Th,  Ce,  Zr,  Y,  La,  Di,  etc.). 

3.  (NH4)2CO4  precipitates  the  carbonate,  readily  soluble  in 

excess,  but  not  so  readily  soluble  in  alkali  carbon- 
ates. 

Thorium.  The  purest  form  of  thorium  found  in  nature  is 
thorite,  ThSiO4 ;  but  it  is  also  found  associated  with  other  rare 
elements  in  gadolinite  (see  Be)  and  in  monazite  sand  (see  also 
Ce),  from  which  nearly  all  commercial  thorium  compounds  are 
derived.  Many  other  rare  minerals  also  contain  it.  Thorium 
compounds  are  of  increasing  interest,  because  of  the  use  of  the 
oxide  in  the  formation  of  gas  mantles  for  the  Welsbach  light. 
The  mantles  are  prepared  by  saturating  a  cotton  form  with 
nitrates  or  bromides  of  the  rare  earths.  The  cotton  web  is 
burned  away,  and  the  somewhat  fragile  shell  of  oxides  which  re- 
mains constitutes  the  mantle  and  is  chiefly  thorium  oxide,  ThO2. 
The  illuminating  power  of  the  mantle  is  said  to  be  at  its  maxi- 
mum when  it  consists  of  99  per  cent  thorium  oxide  and  I  per  cent 
cerium  oxide,  though  the  mixture  usually  contains  oxides  of 
several  other  of  the  rare  metals.  Thorium  salts  are  also  of  great 
interest  because  of  their  radioactivity  and  their  decomposition 
into  thorium  "  emanations."  For  detailed  information,  consult 
works  on  Radioactivity. 

Thorium  minerals  are  usually  decomposed  by  sulphuric  acid, 
and  a  solution  suitable  for  the  characteristic  tests  may  be  obtained 
by  digesting  thorite,  or  monazite  sand,  or  discarded  Welsbach 
mantles,  in  concentrated  H2SO4,  heating  the  mass  to  redness. 
Cool  and  extract  with  cold  water.  The  filtrate  may  be  treated 
with  oxalic  acid  and  the  precipitated  oxalates  dissolved  in  hot 


THE  RARE  METALS  159 

ammonium  oxalate.  The  filtrate  on  acidification  with  dilute 
HC1  will  give  a  precipitate  of  the  oxalate,  and  this  on  ignition 
is  converted  into  almost  pure  thorium  oxide.  This  oxide  dis- 
solved in  concentrated  sulphuric  acid  reacts  as  follows  :  (This 
solution  requires  long  digestion.) 

1.  NH4OH  and  other  bases  precipitate  Th(OH)4  ;   insol- 

uble in  excess  ;  readily  soluble  in  dilute  acid. 

2.  K2CO3  precipitates  the  white  carbonate  ;  soluble  in  ex- 

cess ;  rendered  turbid  by  heating,  and  redissolving 
•   when  cooled. 

3.  Oxalic  acid  and  ammonium  oxalate  react  as  indicated  in 

the  preparation  of  the  solution  (see  above). 

4.  H2O2  gives  a  white   precipitate,  soluble   in   sulphuric 

acid,  the  solution  giving  the  perchromate  test  with 
chromic  acid.     Zirconium  gives  the  same  reaction. 

5.  No  bead  coloration  is  formed  with  thorium  compounds. 

For  fuller  information  concerning  the  analysis  of 
monazite  sand,  see  Zeit.  f.  Angew.  Chem.,  Vol.  14, 
p.  655  ;  1901  ;  also  Ber.,  Vol.  35,  p.  2826;  also  Jour. 
Am.  Chem,  Soc.,  Vol.  24,  p.  901  ;  1902.  For  separa- 
tion of  cerium  and  thorium  and  zirconium,  see  Zeit. 
f.  Anal.  Chem.,  Vol.  36,  p.  676,  and  Vol.  37,  p.  94, 
.  Prakt.  Chem.  (2),  Vol.  66,  p.  59. 


Cerium.  Cerium  occurs  in  monazite  sand,  associated  with 
lanthanum,  neo-  and  praesodymium,  thorium,  and  other  rare 
elements,  as  the  phosphate,  and  as  the  silicate  in  cerite  with 
lanthanum  and  didymium,  and  many  other  rare  minerals.  The 
chief  interest  of  cerium  lies  in  its  relation  to  thorium  in  the 
Welsbach  mantles  and  in  the  use  of  the  oxalates  in  medicine  as 
a  nerve  sedative  in  hysteria  and  pregnancy  and  as  a  remedy  for 
sea  sickness. 

The  element  may  be  extracted  from  its  ores  by  digestion  with 


160  QUALITATIVE  ANALYSIS 

hot  sulphuric  acid  (see  Thorium)  and  a  solution  prepared  as  fol- 
lows: Cool  the  hot  mass  from  which  the  excess  H2SO4has  been  ex- 
pelled, digest  with  cold  water,  and  filter.  The  filtrate  is  treated  with 
H2S  to  remove  second  group  metals,  and  after  filtration  is  treated 
with  oxalic  acid.  The  precipitate  is  digested  with  ammonium 
oxalate.  (See  Thorium.)  The  residual  oxalates  are  ignited  and 
consist  of  cerium,  lanthanum,  and  didymium  oxides,  and  are  dis- 
solved in  concentrated  HC1.  The  solution  is  treated  with  KOH 
in  excess  and  chlorine  passed  into  the  mixture  until  the  precipi- 
tate is  deep  yellow  in  color.  The  mixture  is  now  filtered.  The 
filtrate  contains  the  lanthanum  and  didymium.  See  below. 

The  eerie  oxide  is  dissolved  in  HC1  (i-i)  with  evolution  of 
chlorine,  and  the  clear  solution  of  CeCl3  used  for  the  following 
reactions : 

1.  NH4OH  and  other  bases  precipitate  Ce(OH)3  ;  insoluble 

in  excess. 

2.  H2O2  in  neutral  solution  precipitates  an  orange-yellow 

CeO3,  made  lighter  in  color  by  boiling. 

3.  Bead  test :  Cerium  salts  color  the  oxidized  bead  of  both 

borax  and  microcosmic  salt  orange-yellow  while  hot, 
pale  yellow  when  cold  ;  the  reduced  bead  is  colorless. 

Indium  and  Gallium.  These  elements  are  of  extremely  rare 
occurrence,  mainly  in  zinc  ores,  and  were  discovered  by  means 
of  the  spectroscope.  The  student  is  referred  for  fuller  informa- 
tion regarding  them,  to  Browning's  Introduction  to  the  Rarer 
Elements. 

Lanthanum,  Didymium  {Neo  and  Prceso),  Yttrium,  Scandium, 
Erbium,  etc.  These  elements  are  extremely  rare,  and  together 
with  several  others,  some  of  which  are  of  doubtful  existence, 
are  extracted  from  the  cerium  and  thorium  and  beryllium  min- 
erals. In  case  one  of  these  elements  is  found  in  a  mineral  and 
it  is  desired  to  investigate  for  the  rarer  elements,  the  student  is 


THE  RARE  METALS  l6l 

referred  to  the  system  of  analysis  of  Noyes,  Tech.  Quart.,  Vol. 
16,  pp.  93-131,  and  Jour.  Am.  Chem.  Soc.,  Vol.  29,  p.  137;  1907. 
Also  a  suggested  course  of  procedure  is  given  in  Bottger's  Quali- 
tative Analyse,  2d  Auflage,  p.  208  et  seq. 

Group  IV 

Ammonium  Carbonate  Group 

Rd 

Radium  occurs  in  extremely  minute  quantities  in  the  two 
uranium  ores,  pitchblende  and  carnotite,  from  which  it  was  ex- 
tracted with  almost  infinite  pains  by  Madame  Curie.  It  natu- 
rally has  no  place  in  these  notes,  except  as  it  deserves  mention 
because  of  the  close  resemblance  of  its  chemical  behavior  to 
that  of  barium. 

Group  V 
The  Soluble  Group 

Li,  Rb,  Cs 

Lithium.  This  element  is  found  very  widely  distributed  in 
very  small  amounts  in  tourmaline,  muscovite,  epidote,  and  ortho- 
clase,  which  are  common  silicate  minerals.  It  occurs  also  in 
larger  quantities  in  the  somewhat  rare  lepidolite,  or  lithia  mica, 
triphylite,  petalite,  etc.  Because  of  this  wide  distribution  it 
finds  its  way  into  many  surface  waters  where  its  detection  is 
important.  Traces  of  it  are  also  found  in  many  plant  ashes, 
although  large  amounts  in  soil  seem  to  be  detrimental  to  plant 
growth.  It  is  present  particularly  in  seaweed,  tobacco,  cocoa, 
coffee,  sugar  cane,  and  in  milk,  blood,  and  muscular  tissue.  Its 
salts  find  use  in  medical  preparations,  and  mineral  waters  con- 
taining it  in  weighable  quantities  are  considered  valuable  be- 
cause of  the  solubility  of  lithium  urate.  They  are  therefore 
supposed  to  assist  in  the  removal  of  uric  acid  from  the  tissues 
of  the  body. 


162  QUALITATIVE  ANALYSIS 

Most  of  the  famous  mineral  spring  waters,  as  Durkheim, 
Kissingen,  Ems,  Karlsbad,  Buffalo  Lithia,  and  numerous  others, 
owe  their  curative  value,  real  or  supposed,  to  the  presence  of 
these  salts.  Its  salts 'have  the  general  behavior  of  those  of 
sodium  and  potassium,  but  it  somewhat  resembles  the  alkaline 
earth  metals  in  that  it  forms  a  somewhat  difficultly  soluble  car- 
bonate and  phosphate. 

In  the  regular  course  of  analysis,  it  is  found  along  with 
sodium  and  potassium  in  the  fifth  group  filtrate,  and  if  the  solu- 
tion is  evaporated  to  dryness,  the  dry  chloride  of  lithium  will 
dissolve  in  alcohol  or  in  amyl  alcohol.  The  alcoholic  solution 
may  be  evaporated  and  the  residue  dissolved  in  water.  If 
lithium  is  present,  it  will  give  the  following  reactions : 

1.  Flame    test  :    Lithium    salts    color   the   flame   carmine, 

violet  through  blue  glass  and  invisible  through 
green  glass.  They  also  give  a  very  characteristic 
spectrum,  the  red  line  of  which  is  nearer  the  yellow 
than  that  of  potassium. 

2.  H2PtCl6  forms  no  precipitate. 

3.  Na2HPO4  precipitates  Li3PO4  from  moderately  concen- 

trated solutions,  and  if  the  solution  is  made  slightly 
alkaline  with  NaOH  and  evaporated  to  dryness, 
the  phosphate  remains  almost  quantitatively  insol- 
uble in  water  containing  a  little  ammonia. 

4.  (NH4)2CO3  precipitates  from  the  concentrated  solution 

Li2CO3  (solubility  13  g.  per  liter  at  18°). 

Rubidium  and  Cczsium.  These  elements  occur  widely  dis- 
tributed, but  in  minute  amounts  in  the  minerals  and  mineral 
waters  which  contain  potassium  and  lithium.  Their  chief  claim 
to  interest  lies  in  their  discovery  by  means  of  the  spectroscope 
in  the  hands  of  its  inventors,  Kirchhoff  and  Bunsen. 

They  form   even    more   insoluble  chlorplatinates  than   does 


THE  RARE  METALS  163 

potassium,  and  in  the  regular  course  of  analysis  are  sought  in 
the  potassium  chlorplatinate  precipitate.  After  ignition,  the 
precipitate  is  extracted  by  water  and  the  solution  evaporated  to 
near  dryness.  The  presence  of  rubidium  may  be  recognized  by 
two  dark  red  lines  near  the  potassium  line,  and  two  lines  in  the 
violet.  Caesium  may  be  recognized  by  two  sky-blue  lines. 
Rubidium  colors  the  flame  red  and  caesium  sky-blue. 


APPENDIX 


LIST   OF   APPARATUS 

The  student  will  need  the  following  set  x>f  apparatus  to  carry 
on  the  work  outlined  in  this  manual : 


Asbestos  board,  4x4  inches. 
Beakers,  Nos.  i,  2,  and  3. 
Blue  glass. 

Bottle  (rubber  stopper)  for  NaOH. 
Bottles,  reagent,  glass-stoppered. 
Bunsen  burners,  with  two  feet  of  rub- 
ber tubing. 
Crucible  tongs. 
Evaporating  dishes. 
Filter  papers. 
Flasks,  Erlenmeyer : 

60  c.c. 

100  c.c. 

300  c.c.,  for  H2S  precipitation,  with 

stopper. 
Flasks,  Florence : 

100  c.c.,  for  hydrogen  generator 
(with  rubber  stopper,  funnel  tube, 
and  rubber  connections). 

750  c.c.,  for  wash  bottle,  with  stopper. 

Besides  the  above  apparatus,  the  student  should  supply  him- 
self with  the  following : 


Forceps. 
Funnel  rack. 
Funnels. 
Glass  tubing. 
Platinum  foil. 
Platinum  wire. 
Rubber  tubing. 
Sand  bath. 
Spatula,  bone. 
Stand  with  two  rings. 
Test-tube  brush. 
Test-tube  holder. 
Test-tube  rack. 
Test  tubes,  large. 
Test  tubes,  small. 
Triangles,  plain. 
Watch  glasses. 
Wire  gauzes. 


Apron,  rubber  or  denim. 
Half  sleeve,  pair. 


Matches. 

Notebook. 

Towel. 


165 


1 66  QUALITATIVE  ANALYSIS 

REAGENTS   IN    SOLUTION 

(Needed  in  analysis) 

ACIDS  r 

Acetic,  Sp.  Gr.  1.044:  one  part  glacial  acid  and  two  and 
one  half  parts  of  water. 

Aqua  regia :  one  part  of  nitric  and  three  parts  of  hydrochloric. 

Hydrochloric,  concentrated,  Sp.Gr.  1.20:  36.5  per  cent  HC1. 

Hydrochloric,  dilute  :  one  part  concentrated  acid  and 
four  parts  of  water. 

Hydrochloric,  dilute,  for  separation  of  Co  and  Ni  from  Mn 
and  Zn  :  one  part  concentrated  acid  to  twenty  of  water. 

Hydrogen  sulphide ;  is  used  either  as  a  gas  or  saturated 
solution. 

Nitric,  concentrated,  Sp.  Gr.  1.42  :  68  per  cent  HNO3. 

Nitric,  dilute :  one  part  concentrated  acid  and  four  parts  of 
water. 

Sulphuric,  concentrated  Sp.  Gr.  1.84;  98  per  cent  H2SO4. 

Sulphuric,  dilute :  one  part  concentrated  acid  and  four 
parts  of  water.  (Add  the  acid  to  the  water  with  con- 
stant stirring.) 

Sulphurous  acid  :  used  either  as  a  gas  or  saturated  solution. 
ALKALIES  : 

Ammonium  hydroxide,  concentrated,  Sp.  Gr.  0.90 :  28  per 
cent  NH3. 

Ammonium  hydroxide,  dilute :  one  part  concentrated  am- 
monia and  three  parts  of  water. 

Potassium  hydroxide:  100  grams  of  the  solid  to  one  liter 
of  water. 

Sodium  hydroxide :    100  grams  of  the  solid  to  one  liter  of 

water. 
SALTS  AND  OTHER  REAGENTS  IN  LIQUID  FORM  : 

Ammonium  acetate :  1000  c.c.  of  concentrated  ammo- 
nium hydroxide,  1250  c.c.  of  glacial  acetic  acid,  neu- 


APPENDIX  167 

tralize  with  ammonia  or  acetic  acid;  or  a  saturated 
solution  of  the  salt  in  water. 

Ammonium  sesquicarbonate :  500  grams  of  the  crystals  of 
ammonium  carbonate  in  600  c.c.  of  water  and  80  c.c. 
of  concentrated  ammonia;  after  solution  is  complete, 
make  up  to  a  liter  with  water.  If  the  salt  is  dry,  use 
200  grams  instead  of  500. 

Ammonium  carbonate :  500  grams  of  the  crystals  in  600 
c.c.  of  water ;  after  solution  is  complete,  make  up  to 
a  liter  with  water.  If  the  salt  is  dry,  use  200  grams 
instead  of  500. 

Ammonium  chloride  :  100  grams  of  the  salt  in  a  liter  of  water. 

Ammonium  molybdate :  dilute  100  c.c.  of  ammonia  (Sp. 
Gr.  0.90)  with  150  c.c.  of  water  and  dissolve  in  this 
50  grams  of  molybdic  acid ;  dilute  250  c.c.  of  concen- 
trated nitric  acid  with  500  c.c.  of  water,  and  pour 
slowly  into  the  first  solution  with  constant  stirring ; 
allow  to  stand  in  a  warm  place  for  forty-eight  hours 
and  decant  the  clear  supernatant  liquid  for  use. 

Ammonium  oxalate :  a  saturated  solution  in  water. 

Ammonium  sulphate :  250  grams  of  the  crystallized  salt  in 
a  liter  of  water. 

Ammonium  sulphide,  colorless :  saturate  a  liter  and  a  half 
of  concentrated  ammonium  hydroxide  with  hydrogen 
sulphide;  add  a  liter  of  concentrated  ammonium  hy- 
droxide and  two  liters  of  water. 

Ammonium  sulphide,  yellow :  add  to  the  above  solution 
about  75  grams  of  powdered  roll  sulphur. 

Barium  carbonate :  60  grams  suspended  in  a  liter  of  water. 

Barium  chloride :  60  grams  of  the  crystallized  salt  in  a  liter 
of  water. 

Barium  hydroxide  :  50  grams  of  the  crystallized  salt  in  a 
liter  of  water. 


168  QUALITATIVE  ANALYSIS 

Bromine  water :  50  grams  of  potassium  bromide  in  500  c.c. 
of  water  and  shake  with  10  grams  of  bromine  until 
dissolved. 

Brucine  solution :  0.2  grams  dissolved  in  100  c.c.  concen- 
trated sulphuric  acid. 

Calcium  hydroxide  :  a  saturated  solution  in  water. 

Calcium  sulphate :  a  saturated  solution  in  water. 

Chlorine  water :  water  saturated  with  chlorine  gas.  Keep 
in  a  dark  place  or  in  a  dark  bottle. 

Ether ;  commercial. 

Ethyl  alcohol :  95  per  cent. 

Ferric  chloride :   100  grams  to  a  liter  of  water. 

Ferrous  sulphate :  200  grams  per  liter ;  add  5  c.c.  H2SO4 
and  some  pieces  of  metallic  iron  (tacks) ;  keep  stop- 
pered. 

Hydrogen  peroxide  :  3  per  cent  solution. 

Indigo :  place  five  parts  of  fuming  sulphuric  acid  in  a 
beaker  which  is  immersed  in  water,  and  then  add 
slowly  with  constant  stirring  one  part  of  powdered 
indigo  ;  cover  the  beaker  and  allow  to  stand  forty-eight 
hours,  and  then  pour  into  twenty  times  its  volume  of 
water  ;  stir  thoroughly  and  filter  if  necessary. 

Lead  acetate :   100  grams  of  the  crystals  in  a  liter  of  water. 

Magnesia  mixture :  90  grams  of  magnesium  chloride,  240 
grams  ammonium  chloride,  50  c.c.  concentrated  am- 
monium hydroxide ;  dilute  to  a  liter. 

Methyl  alcohol ;  commercial. 

Mercuric  chloride :  50  grams  to  a  liter  of  water. 

Nessler's  reagent :  7  grams  of  potassium  iodide  are  dis- 
solved in  20  c.c.  of  water,  and  to  this  solution  is  added 
gradually  a  solution  of  3.2  grams  of  mercuric  chloride 
in  60  c.c.  of  water  until  a  permanent  precipitate  is 
formed;  120  c.c.  of  concentrated  potassium  hydroxide 


APPENDIX  169 

solution  is  added  and  allowed  to  stand  until  settled; 
decant  from  the  precipitate  as  needed. 

Phenol  sulphonic  acid  :  dissolve  24  grams  of  phenol  crystals 
in  a  mixture  of  148  c.c.  of  concentrated  sulphuric  acid 
and  12  c.c.  of  water;  keep  in  the  dark. 

Potassium  chromate  :   100  grams  to  a  liter  of  water. 

Potassium  cyanide :  30  grams  to  a  liter  of  water. 

Potassium  ferricyanide  :   10  grams  to  a  liter  of  water. 

Potassium  ferrocyanide :   1 5  grams  to  a  liter  of  water. 

Potassium  fluoaluminate  (reagent  for  sodium) :  treat  an  ex- 
cess of  freshly  precipitated  A1(OH)3  with  concentrated 
HF  in  a  lead  or  platinum  dish.  Allow  to  stand  at 
room  temperature  for  two  days.  Add  an  equal  volume 
of  a  saturated  solution  of  potassium  acetate ;  boil  and 
filter.  Add  an  amount  of  50  per  cent  alcohol  equal  to 
the  volume  of  nitrate.  Filter  again,  if  necessary,  and 
preserve  in  a  glass  container. 

Potassium  iodide :  20  grams  to  a  liter  of  water. 

Potassium  nitrite  :  50  grams  to  a  liter  of  water. 

Potassium  sulphate :  85  grams  to  a  liter  of  water. 

Potassium  sulphocyanate  :  50  grams  to  a  liter  of  water. 

Silver  nitrate  :  25  grams  to  a  liter  of  water. 

Silver  sulphate  :  a  saturated  solution. 

Sodium  acetate :  a  saturated  solution. 

Sodium  ammonium  phosphate  :  70  grams  to  a  liter  of  water. 

Sodium  carbonate :  1 50  grams  of  the  anhydrous  salt  or 
250  grams  of  the  crystallized  salt  to  the  liter. 

Sodium  cobaltic  nitrite :  dissolve  100  grams  of  sodium 
nitrite  in  300  c.c.  of  water,  add  acetic  acid  to  slight 
acid  reaction,  and  then  add  10  grams  of  cobalt  nitrate. 
Allow  to  stand  several  hours  and  filter  if  not  clear. 
The  solution  decomposes  slowly,  therefore  it  is  ad- 
visable to  prepare  only  in  small  quantities. 


170  QUALITATIVE  ANALYSIS 

Stannous  chloride :  heat  an  excess  of  tin  with  concen- 
trated hydrochloric  acid ;  dilute  with  four  volumes 
of  water  and  keep  in  stoppered  bottles  with  an  excess 
of  tin. 

Tartaric  acid  :   100  grams  to  the  liter. 

Zinc  sulphate:  140  grams  of  the  crystallized  salt  to  the 
liter. 

OTHER  SOLUTIONS  USED  IN  PRELIMINARY 
EXPERIMENTS 

Aluminium  sulphate :  30  grams  to  the  liter. 

Antimony   chloride :    25    grams   of    the  salt,    250   c.c.    of 

concentrated    hydrochloric  acid,   and    sufficient  water 

to  make  up  to  a  liter. 
Arsenous  chloride:    35   grams  of   arsenous  oxide,  50  c.c. 

concentrated   hydrochloric   acid,  and  sufficient  water 

to  make  up  to  a  liter. 
Bismuth   chloride :    30    grams    to   a    liter   of    water    and 

enough  hydrochloric  acid  to  dissolve  the  precipitate 

formed. 

Boric  acid  :  a  saturated  solution. 
Cadmium  chloride  :  25  grams  to  a  liter  of  water. 
Calcium  chloride :  25  grams  to  a  liter  of  water. 
Chromium  sulphate :  30  grams  to  a  liter  of  water. 
Cobalt  nitrate  :  30  grams  to  a  liter  of  water. 
Copper  sulphate :  30  grams  to  a  liter  of  water. 
Lead  acetate  (alkaline):  two  parts  lead  acetate,  one  part 

of  ammonium  acetate  in  water ;    make  alkaline  with 

ammonium  hydroxide. 

Magnesium  sulphate  :   100  grams  to  a  liter  of  water. 
Manganese  sulphate :  40  grams  to  a  liter  of  water. 
Mercurous  nitrate :  60  grams  to  a  liter  of  water. 
Potassium  bichromate  :  50  grams  to  a  liter  of  water. 


APPENDIX 


171 


Potassium  bromide  :  30  grams  to  a  liter  of  water. 

Potassium  chloride :  50  grams  to  a  liter  of  water. 

Potassium  chromate  :  50  grams  to  a  liter  of  water. 

Potassium  iodide  :  30  grams  to  a  liter  of  water. 

Sodium  arsenate  :  30  grams  to  a  liter  of  water. 

Sodium  arsenite  :  30  grams  to  a  liter  of  water. 

Sodium  phosphate  :  120  grams  to  a  liter  of  water. 

Stannic  chloride :  1 5  grams  of  stannous  chloride  are 
treated  in  an  evaporating  dish  with  50  c.c.  of  hydro- 
chloric acid  and  5  grams  of  potassium  chlorate,  and 
heated  until  chlorine  is  no  longer  evolved ;  the  solu- 
tion is  then  diluted  to  500  c.c. 

Strontium  chloride  :  30  grams  to  a  liter  of  water. 

REAGENTS   IN    SOLID    FORM 


Aluminium  foil. 

Aluminium  wire. 

Ammonium  chloride. 

Ammonium  nitrate. 

Ammonium  sulphate. 

Ammonium  sulphite. 

Barium  chloride. 

Barium  hydroxide. 

Borax. 

Calcium  carbonate  (marble). 

Copper  foil. 

Copper  wire. 

Cotton,  absorbent. 

Ferrous  sulphate. 

Ferrous  sulphide. 

Iron  filings. 

Iron  wire. 

Lead  acetate. 


Lead  dioxide. 
Oxalic  acid. 
Potassium  carbonate. 
Potassium  chlorate. 
Potassium  cyanide. 
Potassium  nitrate. 
Silver  nitrate. 

Sodium  ammonium  phosphate. 
Sodium  carbonate. 
Sodium  hydroxide. 
Sodium  sulphate. 
Sodium  sulphide. 
Sodium  sulphite. 
Sulphur  (powdered  or  flow- 
ers). 

Tartaric  acid. 
Zinc,  granulated. 
Zinc,  sheet. 


QUALITATIVE  ANALYSIS 


TABLE   OF  ATOMIC   WEIGHTS    (1912) 


ELEMENT 

SYMBOL 

ATOMIC 
WEIGHT 
o  =  16 

ELEMENT 

SYMBOL 

ATOMIC 
WEIGHT 
0=16 

Aluminium  .  .  . 
Antimony  .... 
Ar^on  . 

IP 

A  ' 

27.1 
I2O.2 

30  88 

Molybdenum  .  . 
Neodymium  .  .  . 
Neon  .... 

Mo 

Nd 

Ne 

96.0 

144-3 

2O  2 

Arsenic  

As 

jy.uu 

74..  QO 

Nickel  

.  Ni 

c;8.68 

Barium 

Ba 

I  37  37 

Nitrogen 

N  - 

14  01 

Bismuth  .... 
Boron  

F 

lj/  "Jl 
208.0 
I  I.O 

Osmium  .... 
Oxygen  .... 

Os 

o 

190.9 
16.0 

Bromine  .... 
Cadmium  .... 
Caesium  .... 
Calcium  .... 
Carbon 

Br 

& 

Ca 

c 

79.92 
II2-4 
I32.8I 

40.07 
12  O 

Palladium  .  .  . 
Phosphorus  .  .  . 
Platinum  .... 
Polonium  .... 
Potassium 

Pd 
P 
Pt 

Po 
,  K; 

106.7 
31.04 
195.2 
(?) 

•2Q    IO 

Cerium  .... 

Ce 

I  -dO  25 

Prassodymium  . 

Pr 

jy.iu 

1  40  6 

Chlorine  .... 
Chromium  .... 
Cobalt 

Cl 
Cr 

Co; 

3546 
52.0 
rg  Q7 

Radium  .... 
Rhodium  .... 
Rubidium 

Ra 
Rh 
Rb 

226.4 
102.9 

gt    AC 

Columbium  .  .  . 
Copper  
Dysprosium  .  .  . 
Erbium  

Cb 
Cu 
Dy 
Er 

}«j.y/ 

93.5 

63-57 

162.5 

l67  7 

Ruthenium  .  .  . 
Samarium  .  .  . 
Scandium  .  .  . 
Selenium  . 

Ru 
Sm 
Sc 
Se 

*O'*r3 

IOI.7 
150.4 
44.1 
7Q.2 

Europium  .... 

Eu 

1  12.  0 

Silicon  

Si  - 

28.3 

Fluorine     .... 
Gadolinium     .     .     . 
Gallium      .... 
Germanium     . 
Glucinum   .... 
Gold  .     .     . 

F  - 
Gd 
Ga 
Ge 
Gl 
Au 

19.0 

157-3 
69.9 
72.5 
9.1 

IQ7  2 

Silver  
Sodium  .... 
Strontium  .  .  . 
Sulphur  .... 
Tantalum.  .  .  . 
Tellurium 

Na 
Sr 
S 
Ta 
Te 

107.88 
23.00 
87-63 
32.07 
I8I.5 
127.6* 

Helium  

He 

*y/  •* 

3.  QQ 

Terbium  .... 

Tb 

I  tJQ.2 

Hydrogen  .... 
Indium  . 

H 
In 

j.yy 
1.  008 

114  8 

Thallium  .... 
Thorium 

Tl 
Th 

204.O 
232 

Iodine  

I 

126  92 

Thullium  . 

Tm 

i68.c 

Iridium  

Ir 

-     IQ3.I 

Tin  

Sn 

1  19.0 

Iron  

Fe 

Ayj'* 

CC.8C 

Titanium  .... 

Ti 

48.1 

Krypton  .... 
Lanthanum  .  .  . 
Lead  

Kr 
La 
Pb 

82.92 

139.0 

207.  1 

Tungsten  .... 
Uranium  .... 
Vanadium 

W 
U 
V 

184.0 
238.5 

KI.O 

Lithium 

Li 

6  QA. 

Xenon 

Xe 

I  3O  2 

Lutecium  .... 
Magnesium  .  .  . 
Manganese 

Lu 

Mg 
Mh 

w-y*F 

174.0 
24.32 

r/i   Q-3 

Ytterbium  .  .  . 
Yttrium  .... 
Zinc 

Yb 
Yt 
Zn 

172.0 
89.0 
6c  37 

Mercury  .... 

Hg 

j^f.yj 
200.6 

Zirconium  .  .  . 

Zr 

uio/ 

90.6 

*  Flint,  Am.  four.  ScL,  Vol.  30,  p.  1209,  1910,  gives  atomic  weight  as  124.3. 


APPENDIX 


173 


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APPENDIX  175 


TABLE  OF  METRIC   SYSTEM 

Length                                     Weight  Volume  Notation 

Kilometer    ....  Kilogram     ....  Kiloliter 1000. 

Hectometer      .     .     .  Hectogram  ....  Hectoliter     .     .     .     .  100. 

Decameter  ....  Decagram    ....  Decaliter 10. 

Meter Gram Liter I. 

Decimeter   ....  Decigram     ....  Deciliter .1 

Centimeter  ....  Centigram  ....  Centiliter .01 

Millimeter  ....  Milligram     ....  Milliliter .oof 


METRIC   MEASURES  WITH   ENGLISH   EQUIVALENTS 

MEASURES  OF   LENGTH 

I  Millimeter,  mm.  =   0.03937  inches. 

i  Centimeter,  cm.  =  10  mm.  =  0.3937  inches, 
i  Decimeter,  dm.  =  10  cm.  =  3.9371  inches. 
i  Meter,  m.  =  10  dm.  =  39.3708  inches. 

I  Kilometer,   km.  =  1000  m.  =    0.6214    miles. 

MEASURES  OF  VOLUME 

i  Cubic  centimeter,  c.c.  =        0.06103  cubic  inches. 

(  61.027      cubic  inches,  or, 
i  Cubic  decimeter  (liter),  1.  =  ioooc.c.=  •]  .  0 

<    1.057      U.  S.  quarts. 

i  Cubic  meter  =  1000  1.     =      35.3166    cubic  feet. 

MEASURES   OF   WEIGHT 

I  Milligram,  mg.  =    0.0154  grains, 

i  Gram,         g.  =  1000  mg.  =  15.432    grains. 

I  Kilogram    (kilo)    =  1000  g.     =    2.2046  pounds  (av.). 


ENGLISH   MEASURES  WITH   METRIC   EQUIVALENTS 


I  Inch  =  25.399    millimeters. 
i  Foot  =    0.3048  meters. 
i  Mile  =    1.609    kilometers. 


i  Cubic  inch  =  16.386  cubic  centi- 


meters. 


i  Cubic  foot  =28.315    liters. 


i  Quart          =    0.9463  liters. 

i  Grain  =    0.0648  grams. 

i  Ounce  (av.)  =  28.3496  grains, 
i  Pound  (av.)  =    0.4536  kilograms. 


176  QUALITATIVE  ANALYSIS 

APOTHECARIES'  WEIGHTS  AND  MEASURES  WITH  METRIC  EQUIVALENTS 

i  Grain  =      0.0648  grams. 

i  Scruple  =      1.296    grams. 

i  Drachm  =      3.888    grams. 

i  Ounce  troy  =    31.1035  grams. 

I  Pound  troy  =373.2418  grams. 

i  Minim  =      0.0616  cubic  centimeters. 

i  Fluid  drachm  =      3.6965  cubic  centimeters. 

i  Fluid  ounce  =    29.572    cubic  centimeters. 

i  Pint  =473.11      cubic  centimeters. 


TO   CONVERT  — 

A  —TO—  B  MULTIPLY  BY 

Inches Centimeters 2.54 

Feet Meters 0.305 

Miles Kilometers 1 .609 

Meters .  Inches 39-37 

Kilometers Miles 0.621 

Square  inches Square  centimeters 6.452 

Square  yards Square  meters 0.836 

Square  centimeters      .     .     .  Square  inches °'I55 

Square  meters Square  yards 1.196 

Cubic  inches Cubic  centimeters 16.386 

Cubic  yards Cubic  meters 0.765 

Cubic  centimeters  ....  Cubic  inches 0.061 

Cubic  meters Cubic  yards 1.308 

Fluid  ounces Cubic  centimeters 29-57 

Quarts Liters 0.946 

Cubic  Centimeters      .     .     .  Fluid  ounces 0.034 

Liters.     . Quarts 1-057 

Grains Milligrams 64.799 

Ounces  (av.) Grams 28.35 

Pounds  (av.) Kilograms 0.373 

Grams Grains 15-432 

Kilograms Pounds 2.205 


APPENDIX 


177 


DEGREE   OF   IONIZATION   OF   ACIDS,   BASES,   AND   SALTS 

Unless  otherwise  specified,  the  figures  give  the  per  cent  ionized  in  a  normal 
solution  at  18°,  calculated  from  the  electrical  conductivities : 

ACIDS  PER  CENT  SALTS  PER  CENT 

HNO3 82.0  KC1 75.0 

HN03  (cone.,  62%)   ....     9.6  NH4C1 74.0 

HC1 78-4  NaCl 67.6 

HC1  (cone.,  35%) 13.6  HgCl2 <i.o 

HMnO4  (N/2,  25°)  ....  93.3  KN03 64.0 

HI(N/2,  25°) 90.1  K2SO4 53.0 

HBr  (N/2,  25°) 89.9  K2C03 (49-0) 

HF 7-0  KC103  (N/2) 79.0 

H2S04 51.0  NaHC03 (52.0) 

H2S04  (cone.,  95°/r)  .     ...    0.7  Na2C4H4O6  (N/32, 25°)      .     .(78.0) 

H3PO4  (N/2,  25°)     ....  17.0  ZnSO4 24.0 

HC2H302 0.4  Hg(CN)2 <i.o 

HC2H302  (N/io)      ....     1.3  Ca(S04)  (N/ioo)     .     .     .     .63.0 

H2CO3  (N/io) 0.17 

H2S  (N/io) 0.07 

H3BO3(N/io) o.oi 

HCN  (N/io) o.oi 

BASES  PER  CENT 

LiOH 63.0 

NaOH     ........  73-o 

KOH      ........  77.0 

Ba  (OH)2 69.0 

Ca  (OH)2  (N/64,  25C)  ...  90.0 
Sr  (OH)2  (N/64,  25°)  .  .  .  93.0 
Ba  (OH)2  N/64,  25°)  .  .  .  92.0 
AgOH  (N/I783,  25°)  .  .  .  38.8 

NH4OH 04 

HOH      .     .     .     .  <i  per  10  million 


INDEX 


Acetic  acid,  118,  120. 

Acid  analysis,  101. 

Acid  groups,  103. 

Acids,  bases  and  salts  in  solution,  18. 

Aluminium  reactions,  83. 

Ammonium,  52. 

Ammonium  carbonate  group,  the,  91. 

Ammonium  reactions,  97. 

Ammonium  sulphide  group,  the,  79. 

Amount  of  sample,.  59. 

Amphoterism,  35,  46,  48. 

Analysis  of,  alloys,  127. 

Group  I,  metals,  63. 

Group  II,  subgroup  A,  metals,  69. 

Group  II,  subgroup  B,  metals,  77. 

Group  III,  metals,  87. 

Group  IV,  metals,  94. 

Group  V,  metals,  98. 

Group  I,  acids,  106. 

Group  II,  acids,  no. 

Group  III,  acids,  115. 

Group  IV,  acids,  118. 

Group  V,  acids,  121. 
Antimony  reactions,  75. 
Apparatus,  165. 
Arrhenius  hypothesis,  21. 
Arsenic  acid,  113. 
Arsenic  reactions,  74. 
Arsenous  acid,  113. 
Arsenous  reactions,  72. 
Atomic  weights,  172. 

Balanced  actions,  25. 
Barium  reactions,  92. 
Bases  in  solution,  18. 
Bead  tests,  50,  124,  139. 
Beryllium,  157. 
Bismuth  reactions,  67. 
Boiling  point,  17. 
Boric  acid,  114. 

Cadmium  reactions,  69. 
Caesium,  162. 
Calcium  reactions,  93.  / 
Carbonic  acid,  105. 


Cerium,  159. 

Charcoal  tests,  125. 

Chemical  equilibrium,  law  of,  25,  26. 

Chloric  acid,  118. 

Chromic  acid  reactions,  82,  112. 

Chromium  reactions,  81. 

Closed  tube  test,  the,  123. 

Cobalt  reactions,  86. 

Complex  ions,  32,  38,  45. 

Conductivity  of  acetic  acid,  29. 

Conductors  of  the  first  class,  19. 

Conductors  of  the  second  class,  20. 

Confirmation  of  tests,  59. 

Copper  reactions,  68. 

Decantation,  59. 

Electrolysis,  19. 
Equations,  writing  of,  57. 
Evaporation  of  filtrates,  59. 

Flame  tests,  52,  125,  139. 
Freezing  point  of  solutions,  15. 

Germanium,  152. 
Gold,  145. 
Grouping  of  acids : 

Group  I,  104. 

Group  II,  107. 

Group  III,  112. 

Group  IV,  117. 

Group  V,  119. 
Grouping  of  metals : 

Group  I,  60. 

Group  II,  65. 

Group  III,  79. 

Group  IV,  91. 

Group  V,  95. 
Grouping  of  rarer  elements : 

Group  I,  140. 

Group  II,  145. 

Group  III,  152. 

Group  IV,  161. 

Group  V,  161. 
Gullberg  and  Waage,  law  of,  26. 


179 


i8o 


INDEX 


Hydriodic  acid,  108. 
Hydrobromic  acid,  108. 
Hydrochloric  acid,  107. 
Hydrochloric  acid  group,  the,  60. 
Hydrocyanic  acid,  109. 
Hydroferricyanic  acid,  no. 
Hydroferrocyanic  acid,  109. 
Hydrofluoric  acid,  114. 
Hydrogen  sulphide  group,  the : 

Subgroup  A,  65. 

Subgroup  B,  72. 
Hydrolysis,  42. 
Hydrosulphuric  acid,  104. 
Hypochlorous  acid,  106. 

Indium,  160. 
Introduction,  i. 
Ion  forms  of : 

Group  I,  31. 

Group  II,  32. 

Group  III,  40. 

Group  IV,  51. 

Group  V,  52. 
Ionic  equilibrium,  27. 
lonization  constant,  28. 
lonization  hypothesis,  15. 
lonization  of  acids,  bases,  and  salts,  177. 
Ions,  21. 
Indium,  149. 
Iron  reactions,  80. 

Kahlenberg's  views,  23. 

Lead  reactions,  62,  67. 
Lithium,  161. 

Magnesium,  53. 
Magnesium  reactions,  98. 
Manganese  reactions,  84. 
Mass  action,  law  of,  26. 
Mercury  (ic)  reactions,  65. 
Mercury  (ous)  reactions,  61. 
Microcosmic  bead  tests,  139. 
Molybdenum,  141. 

Nickel  reactions,  86. 
Niobium,  143. 
Nitric  acid,  117. 
Nitrous  acid,  105. 
Non-conductors,  19. 
Notebooks,  60. 

Organic  acids,  120. 
Osmium,  149. 
Osmotic  pressure,  10. 


Ostwald's  dilution  law,  29. 
Oxalic  acid,  115,  119. 
Oxidation,  48. 

Palladium,  148. 
Periodic  system,  173. 
Permanganic  acid,  118. 
Phosphoric  acid,  114. 
Physical  equilibrium,  23. 
Platinum,  147. 
Potassium  reactions,  53,  96. 
Precipitation,  57. 
Preliminary  examination,  122. 
Preparation  of  samples,  126. 
Prussic  acid,  109. 

Qualitative  illustration : 
Group  I,  31. 
Group  II,  32. 
Group  III,  40. 
Group  IV,  51. 
Group  V,  52. 

Radium,  161. 
Raoult's  law,  15,  16. 
Rarer  metals,  138. 
Reagents  in  solid  form,  171. 
Reagents  in  solution,  166. 
Reagents,  test  of,  59. 
Reversible  reactions,  25. 
Rhodium,  150. 
Rubidium,  162. 
Rules  of  solubility,  101,  102. 
Ruthenium,  150. 

Salts  in  solution,  18. 
Selenium,  150. 
Separation  of : 

arsenic,  antimony,  and  tin,  39. 

bismuth,  lead,  and  mercury,  37. 

cobalt  and  nickel,  144. 

copper  and  cadmium,  38. 

iron,  chromium  and  aluminium,  46 

silver,  lead,  and  mercury,  32,  63. 

subgroups  A  and  B,  Group  II,  35. 
Silicic  acid,  104,  115. 
Silver  reactions,  60. 
Sodium  reactions,  53,  96. 
Solubility  product,  30,  41. 
Solubility  table,  174. 
Soluble  group,  the,  95. 
Solute  and  solvent,  8. 
Solution : 

definition,  7. 

hydrates  in,  9. 


INDEX 


181 


Solution : 

kinds,  7. 

properties,  8. 

volume  changes,  8. 
Solvate  theory,  9. 
Stannic  reactions,  77. 
Stannous  reactions,  76. 
Strontium  reactions,  92. 
Sulphuric  acid,  112. 
Sulphurous  acid,  104. 
Supersaturation,  55. 
Systematic  analysis,  122. 

Tables  of  analysis,  130. 
Tables  of  measures,  175. 
Tantalum,  143. 
Tartaric  acid,  115,  120. 
Tellurium,  151. 
Thallium,  140. 
Theory  of  solution,  10. 


Thiosulphuric  acid,  104. 
Thorium,  158. 
Tin  reactions,  76. 
Titanium,  153. 
Tungsten,  142. 

Uranium,  156. 

Valency,  49. 
Vanadium,  152. 
Van't  Hoff's  hypothesis,  13. 
Morse's  modification,  13. 
Van't  HoffV'i,"  16,  22. 
Very  rare  elements,  160. 

Washing  precipitates,  58. 

Zinc  reactions,  85. 
Zirconium,  155. 


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